RP8 EMF of an electrochemical cell

• A salt bridge was used in a cell to measure electrode potential. Explain the function of the salt bridge.

  • The ions in the ionic substance in the salt bridge move through the salt bridge 1
  • To maintain charge balance / complete the circuit
• The standard electrode potential for the reduction of iron(III) ions into iron(II) ions can be measured by connecting a suitable electrode to a standard hydrogen electrode.

  Draw a clearly labelled diagram to show the components and reagents, including their concentrations, in this Fe(III)/Fe(II) electrode

  • Platinum electrode
  • Solution in beaker is a mixture of named soluble iron(II) compound and named soluble iron(III) compound
  • Concentrations of Fe(II) and Fe(III) ions are both 1 mol dm−3
• (a) Identify solution A and give its concentration. State the other essential conditions for the operation of the standard electrode that forms the left-hand side of the cell.

  • Solution A HCl 1.0 moldm^-3
  • Hydrogen at 100kPa
  • 298K
  • B = platinum electrode
  • Inert / unreactive / does not create a potential difference 1
  • Conducts electricity / allows electron flow / conducts / conductor

  
• Explain why the ammeter reading would fall to zero after a time

  • The Fe3+ ions would be used up / reaction completed
• Give one reason, other than cost, why the platinum electrodes are made by coating a porous ceramic material with platinum rather than by using platinum rods.

  • Increases the surface area (so reaction faster)
• State how, if at all, the e.m.f. of this cell will change if the surface area of each platinum electrode is doubled.

  • Unchanged
• The electrode potential of a hydrogen electrode changes when the hydrogen ion concentration is reduced.

  Explain, using Le Chatelier’s principle, why this change occurs and state how the electrode potential of the hydrogen electrode changes. [3 marks]

  Explanation of change:

  • Equilibrium displaced to left (1)
  • to reduce constraint (1)
  • Change in electrode potential: Becomes negative or decreases (1)

  
• State what must be done to maintain the EMF of this fuel cell when in use.

  • (Constantly) add reactants/glucose (and oxygen) OR
  • keep concentration of reactants constant
• Describe an experiment the student does to show that the standard electrode potential for the TiO2+(aq) / Ti(s) electrode is −0.88 V

  (2020 Paper 3)

  • how to prepare the solution of acidified TiO2+(aq)

  Stage 1:

  • Preparing solution (1a) Weigh 7.995 / 8.00 g TiOSO4 - this is by calculating moles (see the question)
  • (1b) Dissolve in / add (allow react with) 0.50 mol dm-3 sulfuric acid
  • (1c) transfer to volumetric flask and make up to the mark
• 2020 Paper 3

  • how to connect the electrodes

  • measurements taken

  • Stage 2: Set up cell (Content can be shown in a labelled diagram)
  • 2a piece of Ti immersed in (1 mol dm−3 acidified) TiO2+(aq) / the solution
  • 2b connect solutions with salt bridge
  • 2c (connect metals through high R) voltmeter
• how the measurements should be used to calculate the standard electrode potential for the TiO2+(aq) / Ti(s) electrode.

  Stage 3: Measurements and calculation

  • (3a) record voltage/potential difference/emf of the cell
  • (3b) Ecell = ERHS – ELHS Ecell = Ecopper – Etitanium
  • (3c) ELHS = ERHS – Ecell OR Ecell should be +1.22 V if Cu on RHS (or −1.22 if Cu electrode on LHS)
• The standard electrode potential, EΘ, for the Fe3+/Fe2+ electrode is +0.77 V

  State the change that needs to be made to the apparatus in the diagram to allow the cell reaction to go to completion.

  • M2 replace voltmeter with lamp/wire/ammeter
  • OR remove voltmeter
• Give one reason why your calculated value is different from the standard electrode potential for Zn2+/Zn electrode.

  • non standard conditions or concentration (of Zn2+)
  • not 1 (mol dm–3 ) or concentration (of Zn2+) less than 1 (mol dm–3 )
  • ALLOW temperature is not 298K