metals

metals

properties
1. physical
  1. shiny
    1. dense
      1. malleable and ductile
        generally high m.p and b.p
        excellent conductors of heat and electricity
2. chemical
  1. form positive ions(cations
    1. metal oxides tend to be basic
      1. react with acid to form hydrogen gas + a salt
metals reactions
1. with acid
  1. If a metal reacts with dilute hydrochloric acid then hydrogen gas and the metal chloride are produced
    1. Metal+HCl(aq) -----> metal chloride+ H2(g)
      1. 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.
2. with water/steam
  1. Reactive metals react with cold water to produce the metal hydroxide and hyrogen gas
    Annotations:
    - 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
    1. Moderately reactive metals react slowly with cold water . However; react rapidly with steam to produce the metal oxide and hydrogen gas.
      Annotations:
      - 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
      1. less reactive(unreactive) metals;
        Pb Cu Ag Au Pt ---> do not react with
      2. Mg is one of the metals used in construction of the Airbus A380
3. with air/oxygen
  1. Many metals react directly with oxygen to form oxides.
    Annotations:
    - Mg Ca Fe Pb
    1. for example;
      1. 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
decomposition of
1. metal nitrates
  1. When nitrates of reactive metals (K,Na) they decompose to produce the metal nitrite and oxygen gas
    1. sodium nitrate ---heat---> sodium nitrite + oxygen
      1. 2NaNO3(s) ---heat---> 2NaNO2(s)+O2
  2. 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
    1. magnesium nitrate ---heat---> magnesium oxide + nitrogen dioxide + oxygen
      1. 2Mg(NO3)2(s) ---heat---> 2MgO(s) + 4NO2(g) + O2(g)
  3. Unreactive metal nitrates(Ag, Pt, Au) produce the metal, nitrogen dioxide and oxygen.
    1. 2AgNO3(s) ---heat---> 2Ag(s) + 2NO2(g) + O2(g)
2. metal carbonates
  1. The carbonates of most reactive metals(K and Na ) thermally stable and require very high temperature to decompose.
    1. Na2CO3(s) ---heat---> no reaction
  2. The carbonates of moderately reactive metals decompose to metal oxide and carbondioxide. Less reactive metal carbonates are too unstable to exist.
    1. CaCO3(s) ---heat---> CaO(s) + CO2(g)
      1. Ag2CO3(s) ---heat---> no reaction
3. metal hydroxides
  1. 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.
    1. Ca(OH)2(s) ---heat---> CaO(s) + H2O(g)
      Annotations:
      - Ca(OH)2(s) : slaked lime CaO(s) : lime
4. metal oxides
  1. 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.
    1. Na2O(s) ---heat---> no reaction
      1. Al2O3(s) ---heat---> no reaction
        2Ag2O(s) ---heat---> 4Ag(s) + O2(g)
uses
1. unreactive
  1. car bodies , coins, pots and pans
  2. uses of aluminium metal
    1. aluminium forms a relatively thick oxide layer on the surface of the metal which prevents further reaction
      1. handles cooking foil
2. reactive
  1. competition reactions
    1. In solid state
      1. 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
  2. in aqueous solution
    1. 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.
      1. 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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	Created by Şule Hilal KARAKUS almost 9 years ago