Oxidation and Reduction +Chemistry+

For a summary of the metals chemical reactions with air/oxygen, acids and oxides/salts (displacement), including word equations and balanced symbol equations, all in the context of the reactivity series just click on its name from this alphabetical order list ... aluminium .. caesium .. calcium .. copper .. francium .. gold .. iron .. lead .. lithium .. magnesium .. platinum .. potassium .. rubidium .. silver .. sodium .. tin .. zinc (but the notes are in reactivity order) and other sub-sections on this page: METAL CORROSION- RUSTING * DISPLACEMENT REACTIONS and OXIDATION - REDUCTION explained

OXIDATION and REDUCTION - REDOX REACTIONS
OXIDATION - definition and examples	REDUCTION - definition and examples
(a) The gain or addition of oxygen by an atom, molecule or ion e.g. ... (1) S ==> SO2 [burning sulphur - oxidised] (2) CH4 ==> CO2 + H2O [burning methane to water and carbon dioxide, C and H gain O] (3) NO ==> NO2 [nitrogen monoxide oxidised to nitrogen dioxide] (4) SO2 ==> SO3 [oxidising the sulphur dioxide to sulphur trioxide in the Contact Process for making sulphuric acid]	(b) The loss or removal of oxygen from a compound etc. e.g. ... (1) CuO ==> Cu [loss of oxygen from copper(II) oxide to form copper atoms] (2) Fe2O3 ==> Fe [iron(III) oxide reduced to iron in blast furnace] (3) NO ==> N2 [nitrogen monoxide reduced to nitrogen, catalytic converter in car exhaust] (4) SO3 ==> SO2 [sulphur trioxide reduced to sulphur dioxide]
(c) The loss or removal of electrons from an atom, ion or molecule e.g. (1) Fe ==> Fe2+ + 2e- [iron atom loses 2 electrons to form the iron(II) ion, start of rusting chemistry] (2) Fe2+ ==> Fe3+ + e- [the iron(II) ion loses 1 electron to form the iron(III) ion] (3) 2Cl- ==> Cl2 + 2e- [the loss of electrons by chloride ions to form chlorine molecules]	(d) The gain or addition of electrons by an atom, ion or molecule e.g. ... (1) Cu2+ + 2e- ==> Cu [the copper(II) ion gains 2 electrons to form neutral copper atoms, electroplating or displacement reaction) (2) Fe3+ + e- ==> Fe2+ [the iron(III) ion gains an electron and is reduced to the iron(II) ion] (3) 2H+ + 2e- ==> H2 [hydrogen ions gain electrons to form neutral hydrogen molecules, electrolysis of acids or metal-acid reaction]
(e) An oxidising agent is the species that gives the oxygen or removes the electrons	(f) A reducing agent is the species that removes the oxygen or acts as the electron donor
REDOX REACTIONS - in a reaction overall, oxidation and reduction must go together
(g) Redox reaction analysis based on the oxygen definitions
(1) copper(II) oxide + hydrogen ==> copper + water CuO(s) + H2(g) ==> Cu(s) + H2O(g) copper oxide reduced to copper, hydrogen is oxidised to water hydrogen is the reducing agent (removes O from CuO) copper oxide is the oxidising agent (donates O to hydrogen) (2) iron(III) oxide + carbon monoxide ==> iron + carbon dioxide Fe2O3(s) + 3CO(g) ==> 2Fe(l) + 3CO2(g) the iron(III) oxide is reduced to iron, the carbon monoxide is oxidised to carbon dioxide CO is the reducing agent (O remover from Fe2O3) the Fe2O3 is the oxidising agent (O donator to CO)] (3) nitrogen monoxide + carbon monoxide ==> nitrogen + carbon dioxide 2NO(g) + 2CO(g) ==> N2(g) + 2CO2(g) nitrogen monoxide is reduced to nitrogen carbon monoxide is oxidised to carbon dioxide CO is the reducing agent and NO is the oxidising agent (4) iron(III) oxide + aluminium ==> aluminium oxide + iron (the Thermit reaction) Fe2O3(s) + 2Al(s) ==> Al2O3(s) + 2Fe(s) iron(III) oxide is reduced and is the oxidising agent aluminium is oxidised and is the reducing agent
(h) Redox reaction analysis based on the electron definitions
(1) magnesium + iron(II) sulphate ==> magnesium sulphate + iron Mg(s) + FeSO4(aq) ==> MgSO4(aq) + Fe(s) this is the 'ordinary molecular' equation for a typical metal displacement reaction, but this does not really show what happens in terms of atoms, ions and electrons, so we use ionic equations like the one shown below. The sulphate ion SO42-(aq) is called a spectator ion, because it doesn't change in the reaction and can be omitted from the ionic equation. No electrons show up in the full equations because electrons lost by x = electrons gained by y!! magnesium + iron(II) ion ==> magnesium ion + iron Mg(s) + Fe2+(aq) ==> Mg2+(aq) + Fe(s) the magnesium atom loses 2 electrons (oxidation) to form the magnesium ion, the iron(II) ion gains 2 electrons (reduced) to form iron atoms. Mg is the reducing agent (electron donor) and the Fe2+ is the oxidising agent (electron remover or acceptor) Displacement reactions involving metals and metal ions are electron transfer reactions. (2) zinc + hydrochloric acid ==> zinc chloride + hydrogen Zn(s) + 2HCl(aq) ==> ZnCl2(aq) + H2(g) the chloride ion Cl- is the spectator ion zinc + hydrogen ion ==> zinc ion + hydrogen Zn(s) + 2H+(aq) ==> Zn2+(aq) + H2(g) Zinc atoms are oxidised to zinc ions by electron loss, so zinc is the reducing agent (electron donor) hydrogen ions are the oxidising agent (gaining the electrons) and are reduced to form hydrogen molecules (3) copper + silver nitrate ==> silver + copper(II) nitrate Cu(s) + 2AgNO3(aq) ==> 2Ag + Cu(NO3)2(aq) the nitrate ion NO3- is the spectator ion copper + silver ion ==> silver + copper(II) ion Cu(s) + 2Ag+(aq) ==> 2Ag(s) + Cu2+(aq) copper atoms are oxidised by the silver ion by electron loss electrons are transferred from the copper atoms to the silver ions, which are reduced the silver ions are the oxidising agent and the copper atoms are the reducing agent (4) iron(II) chloride + chlorine ==> iron(III) chloride (5) halogen (more reactive) + halide salt (of less reactive halogen) ==> halide salt (of more reactive halogen) + halogen (less reactive) X2(aq) + 2KY(aq) ==> 2KX(aq) + Y2(aq) X2(aq) + 2Y-(aq) ==> 2X-(aq) + Y2(aq) where halogen X is more reactive than halogen Y, F > Cl > Br > I X is the oxidising agent (electron acceptor) KY is the reducing agent (electron donor) See GCSE Group 7 The Halogens - displacement reaction notes (6) Electrode reactions in electrolysis are electron transfer redox changes at the negative cathode positive ions are attracted: metal ions are reduced to the metal by electron gain: Mn+ + ne- ==> M n = the numerical charge of the ion and the number of electrons transferred or 2H+(aq) + 2e- ==> H2(g) (for the discharge of hydrogen) at the positive anode negative ions are attracted: negative non-metal ions are oxidised by electron loss e.g. for oxide ions: 2O2- - 4e- ==> O2 or 2O2- ==> O2 + 4e- for hydroxide ion: 4OH- - 4e- ==> O2 + 2H2O or 4OH- ==> O2 + 2H2O + 4e- for halide ions (X = F, Cl, Br, I): 2X- - 2e- ==> X2 or 2X- ==> X2 + 2e-