1.11 Electrode Potentials and Cells

• Standard electrode potential, EƟ, refers to conditions of

  • 298 K
  • 100 kPa
  • 1.00 mol dm−3 solution of ions.
• Cells are used to measure electrode potentials by reference to the standard hydrogen electrode.
• • Solutions are connected by a salt bridge made from filter paper soaked in KNO3
  • This allows flow of ions
  • So balances the charges
• Metal that is easy to oxidise = very negative electrode potential

  Metal that is harder to oxidise has a positive electrode potential
• Standard electrode potentials can be listed as an electrochemical series.
• Calculate EMF of a cell =

  E cell = E reduced - E oxidised
• The conventional representation of cells.

  • Reduced form | Oxidised form || Oxidised form | Reduced form
  • Double vertical lines represent salt bridge
  • Single Vertical separates things in different phases
• Half cells involving solutions of two aqueous ions of the same element

  (ie Fe2+ and Fe3+) - electrode is made of Platinum, so it won't react with the ions
• The importance of the conditions when measuring the electrode potential, E (Nernst equation not required).

  • If standard conditions are maintained, reading on voltmeter when half cell is connected to Standard Hydrogen Electrode = standard electrode potential of the half cell
• State one requirement for the soluble ionic compound used to make the salt bridge [1 mark]

  • Must not react with the electrolyte / ions in solution

  Do not allow ‘must not react’ without further qualification.
• Identify most powerful reducing agent from all species in table [1 mark]

  • Reducing agent is electron donor (so loses electrons)
  • Reduction is gain of electrons, Oxidation is loss of electrons (OIL RIG)
  • Negative, OXIDATION; Positive, Reduction
  • Therefore most negative species in the table
  • Flip the equation around because electrochemical series gives reduction equations - you want species that is OXIDISED
• Suggest why potassium chloride would not be suitable salt for salt bridge in the cell [1]

  • MENTION IONS! = Cl- would react with Cu2+
• Redox reactions take place in electrochemical cells where electrons are transferred from the reducing agent to the oxidising agent indirectly via an external circuit. A potential difference is created that can drive an electric current to do work. Electrochemical cells have very important commercial applications as a portable supply of electricity to power electronic devices such as mobile phones, tablets and laptops. On a larger scale, they can provide energy to power a vehicle.
• • Electrochemical cells can be used as a commercial source of electrical energy.
  • The simplified electrode reactions in a lithium cell
  • Positive electrode: Li+ + CoO2 + e– → Li+ [CoO2]–
  • Negative electrode: Li → Li+ + e–
• • Cells can be non-rechargeable (irreversible), rechargeable or fuel cells.
  • Fuel cells are used to generate an electric current and do not need to be electrically recharged.
• The electrode reactions in an alkaline hydrogen–oxygen fuel cell.

  • Positive electrode 2H2 (g) + 4OH– (aq)  →  4H2O (l) +  4e–      Negative electrode: O2 (g) +  2H2O  +  4e– →  4OH– (aq) 
  • Overall equation: 2H2 + O2 -> 2H2O
• Risks and problems of Hydrogen-oxygen fuel cell

  • Hydrogen is a highly flammable gas
  • Very thick walled cylinders and pipes are needed to store hydrogen which has economic impacts
  • The production of hydrogen is a by-product of the crude oil industry, which means it relies on a non-renewable, finite resource
  • larger containers are needed compared to liquid fuels
• • Some of the problems with lithium ion cells:
  • A global shortage of lithium is likely to make lithium ion cells unsustainable
  • If cells are not recycled but thrown away in landfills, then a huge amount of lithium becomes lost
  • Reports of lithium-ion cell fires have raised concern about the safety of these batteries in electronic devices; it is a reminder to us that lithium is a very reactive element in Group 1 of the periodic table, which is why it has a high electrode potential
• Explain why the ammeter reading would fall to zero after a time

  • The Fe3+ ions would be used up / reaction completed
  • Answer must relate to reactants in 4(e)(i) equation if given
  • Allow reactant / reactants (ions!) used up
  • Do not allow concentration of Fe3+ decreases
  • Allow concentration of Fe3+ falls to zero
• State if EMF of the cell will change if surface area of each platinum electrode is added

  • NO CHANGE
• In the external circuit of the cell, electrons flow through the ammeter from right to left [2 marks]

  • USING THE DIAGRAM, I can see that concentration of Cu2+ (1 M) is much greater than concentration of Cu2+ (0.4M) (MP1)
  • So reaction of Cu2+ will occur in preference in the left hand electrode (E LHS > E RHS)
  • Reduction at left hand electrode and oxidation at right hand electrode (Think NOPR)
• Identify the material from which electrode is made. Give two reasons why this material is suitable for its purpose

  • Platinum
  • Inert
  • Conducts ELECTRICITY / allows electrons to flow
• Identify solution (salt bridge) which could be used in C. Give two reasons why it is suitable for its purpose [2 marks]

  • KNO3 (aq)
  • K+ does not react with electrode
  • Allows IONS to flow
  • Does not react with solution in electrode
• Explain why a fuel cell does not need to be recharged

  • Hydrogen / the fuel / reactants supplied continuously / fed in
• Use these half-equations to explain how an electric current can be generated

  • Hydrogen (electrode) produces electrons
  • Oxygen (electrode) accepts electrons

  OR electrons flow to the oxygen electrode
• To provide energy for a vehicle, hydrogen can be used either in a fuel cell or in an internal combustion engine. Suggest the main advantage of using hydrogen in a fuel cell rather than in an internal combustion engine.

  • In the fuel cell, a greater proportion of the energy available from the hydrogen-oxygen reaction is converted into useful energy
  • Allow less energy wasted / more efficient
• Identify one major hazard associated with the use of a hydrogen–oxygen fuel cell in a vehicle.

  • Hydrogen is flammable / H+ corrosive / OH– corrosive / hydrogen explosive
• Solar cells generate an electric current from sunlight. These cells are often used to provide electrical energy for illuminated road signs. Explain why rechargeable cells are connected to these solar cells.

  • Solar cells do not supply electrical energy all the time
  • Rechargeable cells can store electrical energy for use when the solar cells are not working
• Suggest one reason why many waste disposal centres contain a separate section for cells and batteries.

  • Prevent pollution of the environment by toxic or dangerous substances / recycling of valuable components
• Suggest a type of cell that behaves like Cell X

  • fuel cell
  • BECAUSE reagents / fuel supplied continuously
  • concentrations (of reagents) remain constant

  
• Positive electrode in alkaline hydrogen-oxygen fuel cell

  • HYDROGEN reacts with hydroxide ions (alkaline)
  • To form water and electrons
  • H2 + 2OH- -> 2H2O + 2e-
• Negative electrode in alkaline hydrogen-oxygen fuel cell

  • OXYGEN (fed in) and WATER AND ELECTRONS (from positive electrode reactions)
  • reacts to form HYDROXIDE IONS (a cycle occurs)
  • O2 + 2H2O + 4e- -> 4OH-