metals

metals

1. properties

   • physical
     • shiny
       • dense
         • malleable and ductile
           • generally high m.p and b.p
             • excellent conductors of heat and electricity
   • chemical
     • form positive ions(cations
       • metal oxides tend to be basic
         • react with acid to form hydrogen gas + a salt

2. metals reactions

   • with acid
     • If a metal reacts with dilute hydrochloric acid then hydrogen gas and the metal chloride are produced
       • Metal+HCl(aq) -----> metal chloride+ H2(g)
         • example
           • magnesium+hydrochloric acid ------> magnesium chloride+hydrogen
             • Mg(s)+2HCl(aq)------>MgCl2(aq)+H2(g)
               • If a similar reactions are carried out using other metals with acid, an order of reactivity can be produced by measuring the rate of evolution of hydrogen. This is known as a reactivity series.
                 • The most reactive metal is the one that has the highest tendency to lose outer electrons to from a positive metal.
                   • The higher evolution rate, the more reactive metal.
   • with water/steam
     • Reactive metals react with cold water to produce the metal hydroxide and hyrogen gas
       •  1.  2K(s)+ 2H2O(l) --->2 KOH(aq)+H2(g)  2. 2Na(s)+ 2H2O(l) ---> 2NaOH(aq)+H2(g)  3. Ca(s)+H2O(l) ----> Ca(OH)2(aq)+H2(g) Rate of evolution of hydrogen; 1>2>3
       • Moderately reactive metals react slowly with cold water . However; react rapidly with steam to produce the metal oxide and hydrogen gas.
         • 1. Mg(s)+H2O(g) ----> MgO(s)+H2(g) 2. Zn(s)+H2O(g) -----> ZnO(s)+H2(g) 3. 2Fe(s)+3H2O(g) --> Fe2O3(s)+3H2 Rate of evolution of hydrogen; 1>2>3
         • less reactive(unreactive) metals;
           • Pb Cu Ag Au Pt ---> do not react with
         • Mg is one of the metals used in construction of the Airbus A380
   • with air/oxygen
     • Many metals react directly with oxygen to form oxides.
       • Mg Ca Fe Pb
       • for example;
         • magnesium burns brightly in oxygen to form the white powder magnesium oxide
           • magnesium+oxygen ---> magnesium oxide
             • 2Mg(s)+O2(g)----> 2MgO(s)
               • Ag ,Au and Pt do not react with oxygen

3. decomposition of

   • metal nitrates
     • When nitrates of reactive metals (K,Na) they decompose to produce the metal nitrite and oxygen gas
       • sodium nitrate ---heat---> sodium nitrite + oxygen
         • 2NaNO3(s) ---heat---> 2NaNO2(s)+O2
     • Nitrates of moderately reactive metals (Ca, Mg, Al, Zn, Fe, Pb, Cu) produce brown fumes of nitrogen dioxide gas when heated, as well as the metal oxides and oxygen gas
       • magnesium nitrate ---heat---> magnesium oxide + nitrogen dioxide + oxygen
         • 2Mg(NO3)2(s) ---heat---> 2MgO(s) + 4NO2(g) + O2(g)
     • Unreactive metal nitrates(Ag, Pt, Au) produce the metal, nitrogen dioxide and oxygen.
       • 2AgNO3(s) ---heat---> 2Ag(s) + 2NO2(g) + O2(g)
   • metal carbonates
     • The carbonates of most reactive metals(K and Na ) thermally stable and require very high temperature to decompose.
       • Na2CO3(s) ---heat---> no reaction
     • The carbonates of moderately reactive metals decompose to metal oxide and carbondioxide. Less reactive metal carbonates are too unstable to exist.
       • CaCO3(s) ---heat---> CaO(s) + CO2(g)
         • Ag2CO3(s) ---heat---> no reaction
   • metal hydroxides
     • Hydroxide of reactive metals show no decomposition when they are heated. The hydroxide of moderately rractive metals do decompose to produce the metal oxide and water.
       • Ca(OH)2(s) ---heat---> CaO(s) + H2O(g)
         • Ca(OH)2(s) : slaked lime CaO(s) : lime
   • metal oxides
     • It is too difficult to decompose the oxides of reactive and moderately reactive metals. It is possible to thermally decompose some oxides of less reactive metals.
       • Na2O(s) ---heat---> no reaction
         • Al2O3(s) ---heat---> no reaction
           • 2Ag2O(s) ---heat---> 4Ag(s) + O2(g)

4. uses

   • unreactive
     • car bodies , coins, pots and pans
     • uses of aluminium metal
       • aluminium forms a relatively thick oxide layer on the surface of the metal which prevents further reaction
         • handles cooking foil
   • reactive
     • competition reactions
       • In solid state
         • A more reactive metal has a greater tendency to form a metal ion by losing electrons than a less reactive metal does. If a more reactive metal is heated with the oxide of less reactive metal, then it will remve the oxygen from it.
           • example:
             • Iron(ııı) oxide is mixed with aluminium and the mixture heated by using magnesium fuse, a very violent reaction occurs as the competititon between the aluminium and the iron for the oxygen take place
               • Fe2O3(s) + 2Al(s)---heat---> Al2O3(s) + 2Fe(s)
                 • 2Fe^3+(s) + 2Al(s) 2Al^3+(s) + 2Fe(s)
                   • This reaction is exothermic and redox reaction. This particular reaction is known as the Thermit reaction.
                     • Since large amounts of heat are given out and the iron is formed in molten state, this reaction is used to weld together damaged railway lines.
                       • Chromium and titanium are prepared from their oxides using this type of competition reaction.
                         • Carbon, non-metal, is able to reduce metal oxides below it in the series.
                           • Fe2O3 (s) + 3CO(g)---heat---> 3CO2(g) + 2Fe(s)
                             • PbO(s) + CO(g) ---heat---> CO2(g) + Pb(s)
                 • or
     • in aqueous solution
       • A more reactive metal will displace less reactive metal from a solution of its salt. This type of reaction is known as a displacement reaction.
         • example
           • When a piece of zinc metal is left to stand in a solution of copper(ıı) sulphate. The copper(ıı) sulphate slowly loses its blue colour as the zinc continiues to displace the copper from the solution and eventually becomes colourless zinc sulphate.
             • Zn(s) + Cu(NO3)2(aq) ---heat---> Zn(NO3)2(aq) + Cu(s)

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