redox and electrode potentials

• Redox


  Reduction-oxidation
• Reducing agents


  Electron donors
• Oxidising agents


  Electron acceptors
• Oxidation


  Process of electron loss
• Oxidation


  Involves an increase in oxidation number
• Reduction


  Process of electron gain
• Reduction


  Involves a decrease in oxidation number
• Redox equations and half equations

  1. Br2 (aq) + 2I- (aq) → I2 (aq) + 2 Br- (aq)

  2. Br2 (aq) + 2e- → 2 Br- (aq)

  3. 2I- (aq) → I2 (aq) + 2 e-
• I has oxidised as it has lost electrons

  Br has reduced as it has gained electrons
• Oxidising agent


  Species that causes another element to oxidise, it is itself reduced in the reaction
• Reducing agent


  Species that causes another element to reduce, it is itself oxidised in the reaction
• Reduction half equation


  Shows the parts of a chemical equation involved in reduction, with electrons on the left
• Oxidation half equation


  Shows the parts of a chemical equation involved in oxidation, with electrons on the right
• Balancing redox equations


  Work out oxidation numbers

  2. Add electrons equal to the change in oxidation number

  3. Check charges balance
• Balancing half equations with O and H

  Balance O with H2O

  2. Balance H with H+

  3. Check charges balance
• Combining half equations


  Multiply half equations to get equal electrons

  2. Add half equations and cancel electrons
• Manganate redox titration is a common exercise, with a significant colour change from reactant to product
• Manganate titration


  Only use dilute sulfuric acid, other acids can interfere
• The purple colour of manganate can make it difficult to see the bottom of the meniscus in the burette
• Thiosulfate redox titration


  Starch indicator is added near the end point to emphasise the colour change
• Manganate titrations with other substances

  H2O2, ethanedioate, iron(II) ethanedioate
• The reaction between MnO4- and C2O4(2-) is slow to begin with, so the conical flask can be heated to 60°C to speed it up
• Manganate titration calculation


  Find moles of KMnO4

  2. Use balanced equation to find moles of reactant

  3. Find mass of reactant

  4. Calculate percentage
• Electrochemical cell


  Has two half-cells connected by a salt bridge
• Salt bridge


  Allows free movement of ions to complete the circuit
• Find moles of KMnO4

  1. moles = conc x vol

  2. 0.02 x 9.8/1000

  3. =1.96x10-4 mol
• Using balanced equation find moles Fe2+ in 10cm3

  1. moles of KMnO4 x 5

  2. =9.8x10-4 mol
• Find moles Fe2+ in 100cm3

  1. 9.8x10-4 mol x 10

  2. =9.8x10-3 mol
• Find mass of Fe in 9.8x10-3 mol

  1. mass= moles x RAM

  2. =9.8x10-3 x 55.8

  3. =0.547g
• Find % mass


  1. %mass = 0.547/2.41 x100

  2. =22.6%
• Electrode Potentials


  Electrochemical cells
• Electrochemical cell


  • A cell has two half–cells
  • The two half cells have to be connected with a salt bridge
  • Simple half cells will consist of a metal (acts an electrode) and a solution of a compound containing that metal
  • These two half cells will produce a small voltage if connected into a circuit
• Salt bridge


  • Used to connect up the circuit, the free moving ions conduct the charge
  • Usually made from a piece of filter paper (or material) soaked in a salt solution, usually potassium nitrate
  • The salt should be unreactive with the electrodes and electrode solutions
• Why does a voltage form?

  • The zinc half-cell has more of a tendency to oxidise to the Zn2+ ion and release electrons than the copper half-cell
  • More electrons will therefore build up on the zinc electrode than the copper electrode
  • A potential difference is created between the two electrodes
• Why use a high resistance voltmeter?

  • The voltmeter needs to be of very high resistance to stop the current from flowing in the circuit
  • In this state it is possible to measure the maximum possible potential difference (E)
  • The reactions will not be occurring because the very high resistance voltmeter stops the current from flowing
• What happens if current is allowed to flow?

  • The voltage will fall to zero as the reactants are used up
  • The most positive electrode will always undergo reduction
  • The most negative electrode will always undergo oxidation
• Measuring the electrode potential of a cell

  • It is not possible to measure the absolute potential of a half electrode on its own
  • It has to be connected to another half-cell of known potential, and the potential difference between the two half-cells measured
  • By convention we can assign a relative potential to each electrode by linking it to a reference electrode (hydrogen electrode), which is given a potential of zero Volts
• The standard hydrogen electrode


  • The potential of all electrodes are measured by comparing their potential to that of the standard hydrogen electrode
  • The standard hydrogen electrode (SHE) is assigned the potential of 0 volts
  • The hydrogen electrode equilibrium is: H2 (g) 2H+ (aq) + 2e-
  • To make the electrode a standard reference electrode some conditions apply: 1. Hydrogen gas at pressure of 100kPa, 2. Solution containing the hydrogen ion at 1 mol dm-3, 3. Temperature at 298K
• Secondary standards


  • The standard hydrogen electrode is difficult to use, so often a different standard is used which is easier to use
  • These other standards are themselves calibrated against the SHE
  • The common ones are: silver / silver chloride, calomel electrode
• Standard Electrode Potentials


  • The standard conditions are: all ion solutions at 1 mol dm-3, temperature 298K, gases at 100kPa pressure, no current flowing
  • When an electrode system is connected to the hydrogen electrode system, and standard conditions apply the potential difference measured is called the standard electrode potential