Redox and electrode potentials

• Redox half equations:

  • Shows the movement of electrons in redox reactions
  • Fe2+ —-> Fe3+ + e- oxidation
  • Fe3+ + e- —-> Fe2+
• • MnO4- —-> Mn2+
  • Balance the electrons. MnO4- + 5e- —> Mn2+
  • Add water to balance the ”O”. MnO4- + 5e- —> Mn2+ + 4H2O
  • Add H+ to balance the water. MnO4- + 5e- + 8H+ —> Mn2+ + 4H2O
• Fe2+ —> Fe3+ + e-

  5Fe2+ —> 5Fe3+ + 5e-

  8H+ + MnO4- + 5e- —> Mn2+ + 4H2O

  8H+ + MnO4- + 5Fe2+ —> 5Fe3+ + 4H2O + Mn2+
• • The redox reaction between iron (II) ions, Fe2+ (aq), and manganate (VII) ions, MnO4- in acid connditions is used as a basis for a redox titration
  • Fe2+ is oxidise to Fe3+
  • MnO4- is reduced to Mn2+
  • The solution containing MnO4- ions is purple & is decolourised by Fe2+ (aq) ions o form a colourless solution containing Mn2+ (aq) ions
• • When a solution of Fe3+ (aq) reacts with iodide ions, I- (aq), the orange-brown Fe3+ (aq) ions are reduced to pale green Fe2+ (aq) ions
  • Fe3+ is reduced to Fe2+
  • I- is oxidised to I2
• • Aqueous dichromate (VI) ions, Cr2O7 2- (aq), have an orange colour & aqueous chromium (III) ions, Cr3+ (aq) have a green colour
  • Acidified Cr2O7 2- (aq) ions can be reduced to Cr3+ (aq) ions by the addition of zinc
  • Cr2O7 2- (aq) + 14H+ (aq) + 3Zn (s) —> 2Cr3+ (aq) + 7H2O (l) + 3Zn2+ (aq)
  • With an excess o zinc, chromium (III) ions are reduced further to chromium (II), which is a pale blue colour
  • Zn(s) + 2Cr3+ (aq) —> Zn 2+ (aq) + 2Cr 3+ (aq)
• • Hot alkaline hydrogen peroxide, H2O2, is a powerful oxidising agent & can be used to oxidise chromium (III) in Cr3+ to chromium (VI) in CrO42-
  • 3H2O2 + 2Cr3+ + 10OH- —> CrO42- + 8H2O
  • Chromium is oxidised from +3 in Cr 3+ to +6 in CrO42-
  • Oxygen is reduced fron -1 in H2O2 to -2 in CrO42-
• • When aqueous copper (II) ions, Cu2+ reacts with excess iodine ions, I- (aq) a redox reaction occurs
  • I- is oxidised to brown iodine I2
  • Cu2+ is reduced to Cu+
  • The Cu+ forms a white precipitate of copper (I) iodide
  • 2Cu2+ (aq) + 4I- (aq) —> 2CuI (aq) + I2 (s)
  • Pale blue —> white precipitate + brown
• • When solid copper (I) oxide Cu2O reacts with hot dilute sulfuric acid, a brown precipitate of copper is formed together with a blue solution of copper(II) Sulfate
  • In this reaction, copper (I) ions Cu+ have been simultaneously oxidised & reduced
  • As the same element has been reduced & oxidised, this reaction is disproportionation
  • Cu2O (s) + H2SO4 (aq) —> Cu (s) + CuSO4 (aq) + H2O (l)
• • When heated with hydroxide ions, NH4+ reacts to produce ammonia gas NH3
  • NH4+ (aq) + OH- (aq) —> NH3(g) + H2O (l)
  • To test for the ammonium ion, aqueous sodium hydroxide, NaOH (aq) is heated gently with the solution being analysed
  • If ammonia is evolved, damp red pH indicator paper will turn blue, confirming the presence of NH4+ ions
• Reducing & oxidising agents:

  • Reduction - gain of electrons or decrease in oxidation number
  • Oxidation - loss of electrons or increase in oxidation number
  • The oxidising agent takes electrons from the species being oxidised. The oxidising agent contains the species that is reduced
  • The reducing agent adds electrons to the species being reduced. The reducing agent contains the species that is oxidising
• Manganate ( VII) titrations can be used for the analysis of many different reducing agents:

  • Iron (II) ions, Fe2+ (aq)
  • ethanedioic acid, (COOH)2 (aq)
• Non- familiar redox titrations:

  • Manganate (VII) titrations can be used to analyse reducing agents that reduce MnO4- to Mn2+
  • KMnO4 can be replaced with other oxidising agents, the commonest used being acidified dichromate (VI), H+/ CR2O7 2-
• What method can determine the copper content of copper (II) salt or alloys?

  Iodine/thiosulfate titrations
• What ions are produced by dissolving copper (II) salts in water?

  Cu$^{2+}$ (aq) ions
• How can insoluble copper (III) compounds be converted to Cu$^{2+}$ ions?


  By reacting with acids
• What is the process for determining copper content in copper alloys?

  React and dissolve in concentrated nitric acid
• What is the reaction equation for copper solid to copper ions?

  Cu (s) --> Cu$^{2+}$ (aq)
• What happens when Cu$^{2+}$ ions react with I$^{-}$ ions?


  They form iodine and copper(I) iodide
• What is the appearance of the mixture formed from Cu$^{2+}$ ions and I$^{-}$ ions?


  Brown colour
• What is the balanced chemical equation for the reaction of Cu$^{2+}$ ions with I$^{-}$ ions?


  2Cu$^{2+}$ (aq) + 4I$^{-}$ (aq) --> 2CuI (s) + I$_{2}$ (aq)
• What is done with the iodine in the brown mixture?

  It is titrated with sodium thiosulfate
• What is the balanced equation for the reaction of iodine with sodium thiosulfate?

  2S$_{2}O₃$^2- (aq) + I$_{2}$ (aq) --> 2I$^{-}$ + S$_{4}O₆$^2- (aq)
• How many moles of Cu$^{2+}$ produce one mole of iodine?


  2 moles of Cu$^{2+}$
• What is the equivalence of 1 mole of Cu$^{2+}$ in terms of sodium thiosulfate?


  1 mole of S$_{2}O₃$^2-
• • Standard electrode potentials are measured using an standard hydrogen electrode (S.H.E)
• Voltaic cell --> converts chemical energy into electrical energy
• Half cell:

  • Contains the chemical species present in a redox half-equation
  • A voltaic cell can be made by connecting together 2 different half-cells, which then allows electrons to flow
  • In the cell, the chemicals in the 2 half-cells must be kept apart - of allowed to mix, electrons would flow in an uncontrollable way & heat energy would be released rather than electrical energy
• Metal / metal ion half-cells:

  • At the phase boundary where the metal is in contact with its ions, an equilibrium will be set up
  • By convention, the equilibrium in a half-cell is written so that the forward reaction shows reduction & reverse reaction shows oxidation
  • In an isolated half-cell, there is no net transfer of electrons either into or out of the metal
  • When 2 half-cells are connected, the direction of electrons flow depends upon the relative tendency of each electrode to release electrons
• Ion/ ion half-cell:

  • An ion/ ion half-cell contains ions of the same element in different oxidation states. E.g: a half cell can be made containing a mixture of aqueous iron (II) and iron (III) ions
  • Fe 3+ (aq) + e- <==> Fe2+ (aq)
  • In this type of half-cell there is no metal to transport electrons either into or out of the half-cell, so an inert metal electrode made out of platinum is used
• Electrode potentials:

  • In a cell with 2 metal/metal ion half-cells connected, the more reactive metal releases electrons more readily & is oxidised
  • In operating cell:
  • the electrode with more reactive metal loses electrons & is oxidised - this is the negative electrode
  • The electrode with the less reactive metal gains electrons & is reduced - this is the positive electrode
• Standard electrode potential:

  • The tendency to be reduced & gain electrons is measured as a standard electrode potential
  • The standard chosen is a half-cell containing hydrogen gas & a solution containing H+ (aq) ions
  • An inert platinum electrode is used to allow electrons into & out of the half-cell
  • The standard conditions used are:-
  • solutions have a concentration of exactly 1 moldm^-3
  • the temperature is 298K (25 degrees celsius)
  • the pressure is 100KPa (1 bar)
  • The standard electrode potential of a standard hydrogen electrode is exactly 0V
• Measuring a standard electrode potential:

  • To measure a standard electrode potential, the half-cell is connected to a standard hydrogen electrode
  • the 2 electrodes are connected by a wire to allow a controlled flow of electrons
  • the 2 solutions are connected with a salt bridge which allows ions to flow
  • the salt bridge typically contains a concentrated solution of an electrolyte that doesn't react with either solution
  • An example of a salt bridge is a strip of filter paper soaked in aqueous potassium nitrate KNO3 (aq)
• The more negative the electrode potential value:

  • The greater the tendency to lose electrons & undergo oxidation
  • The less the tendency to gain electrons & undergo reduction
• The more positive the electrode potential value:

  • The greater the tendency to gain electrons & undergo reduction
  • The less the tendency to lose electrons & undergo oxidation
• • Metals tend to have negative electrode potential values & lose electrons. Non- metals tend to have positive electrode potential values & gain electrons
  • The more negative the electrode potential value, the greater the reactivity of a metal in losing electrons
  • The more positive the electrode potential value the greater the reactivity of a non-metal in gaining electrons
• • Standard electrode potentials essentially quantify the tendency of redox systems to gain or lose electrons
  • A standard cell potential can be calculated directly from standard electrode potentials
  • Electrode cell = most positive electrode potential - negative one
• • An oxidising agent takes electrons away from the species being oxidised. So oxidising agents are reduced & are on the left
  • A reducing agent adds electrons to the species being reduced
  • The strongest reducing agent is at the top on the right of table
  • The strongest oxidising agent is at the bottom on the table
• • The redox system with the more positive electrode value will react from left to right & gain electrons
  • The redox system with the less positive electrode value will react from right to left & lose electrons
• What do electrode potentials indicate about a reaction?

  They indicate thermodynamic feasibility