Commercial Applications of Electrochemical Cells (Oxford AQA International A Level Chemistry)

Commercial Cells

• Electrochemical cells can be used as a commercial source of electrical energy

• Cells can be:

  • Non-rechargeable (irreversible) cells

  • Rechargeable cells

  • Fuel cells

• The type of cells used in commercial applications depend on the:

  • Required voltage

  • Required current

  • Size

  • Cost

• Although it is commonly used incorrectly, the term battery should be used to refer to a collection of cells

  • A car battery is correct, because it is a collection of six cells joined together

Non-rechargeable cells

The Daniell cell

• The Daniell cell was one of the earliest electrochemical cells and consisted of a simple metal-metal ion system

• It was invented by British chemist John Daniell in 1836

• The Daniell cell consists of:

  • A zinc rod immersed in a solution of zinc sulfate

  • A copper cylinder filled with copper sulfate solution

  • A porous pot that separates the copper sulfate from the zinc sulfate

• The zinc acts as the negative electrode and the copper is the positive electrode

• The half-cell reactions are

Zn (s) → Zn²⁺ (aq) + 2e^-           E^ꝋ = -0.76 V 

Cu²⁺ (aq) + 2e^- → Cu (s)              E^ꝋ = +0.34 V

• The cell generates an electromotive force of:

  • E_cell^ꝋ = E_reduction^ꝋ - E_oxidation^ꝋ

  • E_cell^ꝋ = (+0.34) - (-0.76) = +1.10 V

• The overall reaction is:

Zn (s) + CuSO₄ (aq) → Cu (s) + ZnSO₄ (aq)

• However, the cell is impractical to use as a portable device because of the hazardous liquids in the cell

Zinc-carbon cells

• Zinc-carbon cells are the most common type of non-rechargeable cells

• Zinc-carbon cells consist of:

  • A zinc casing

  • A paste of ammonium chloride which acts as an electrolyte as well as the positive electrode

  • A carbon rod which acts as an electron carrier in the cell

• The zinc acts as the negative electrode and the ammonium chloride is the positive electrode

• The half-cell reactions are

Zn (s) →   Zn²⁺ (aq)  +  2e^-                               E^ꝋ = -0.76 V 

2NH₄⁺ (aq) + 2e^- → 2NH₃ (g) + H₂ (g)              E^ꝋ = +0.74 V

• The cell generates an electromotive force of:

  • E_cell^ꝋ = E_reduction^ꝋ - E_oxidation^ꝋ

  • E_cell^ꝋ = (+0.74) - (-0.76) = +1.50 V

• The overall reaction is:

2NH₄⁺ (aq) + Zn (s)  → 2NH₃ (g) + H₂ (g)  + Zn²⁺ (aq)

• One disadvantage is that as the cell discharges, the zinc casing eventually wears away and the corrosive contents of the electrolyte paste can leak out

Rechargeable Cells

• Rechargeable cells employ chemical reactions which can be reversed by applying a voltage greater than the cell voltage, causing electrons to push in the opposite direction

• There are many types of rechargeable cells, but common ones include:

  • Lead-acid batteries

  • NiCad cells

  • Lithium cells (covered in more detail in the next section)

Lead-acid batteries

• Lead-acid batteries consist of six cells joined together in series

• The cells use lead metal as the negative electrode and and lead(IV) oxide as the positive electrode

• The electrolyte is sulfuric acid

• The half-cell reactions are:

Pb (s) +  SO₄^2- (aq)  ⇌   PbSO₄ (s)  +  2e^-                                                 E^ꝋ = -0.36 V 

PbO₂ (s) +  4H⁺ (aq) +  SO₄^2- (aq) +  2e^- ⇌  PbSO₄ (s)  + 2H₂O (l)         E^ꝋ = +1.70 V

• Each cell generates an electromotive force of:

  • E_cell^ꝋ = E_reduction^ꝋ - E_oxidation^ꝋ

  • E_cell^ꝋ = (+1.70) - (-0.36) = +2.06 V

• The overall reaction is:

PbO₂ (s) + 4H⁺ (aq) + 2SO₄^2- (aq) + Pb (s) → 2PbSO₄ (s) + 2H₂O (l)

• In a commercial car battery, the six cells in series give a combined voltage of about 12 V

• When the car is in motion, the generator provides a push of electrons that reverses the reaction and regenerates lead and lead(IV) oxide

• Lead-acid batteries are designed to produce a high current for a short period of time, hence their use in powering a starter motor in car engines

• The disadvantage of lead-acid batteries is that:

  • They are very heavy

  • They contain toxic lead and lead(IV) oxide

  • The sulfuric acid electrolyte is very corrosive

• This presents challenges of disposal when lead-acid batteries come to the end of their useful life

NiCad cells

• NiCad stands for nickel-cadmium and these cells are available in many standard sizes and voltages so they can replace almost any application of traditional zinc-carbon cells

• Although they are more expensive cells, the fact they can be recharged hundreds of times means they are commercially viable

• The positive electrode consists of cadmium and the negative electrode is made of a nickel(II) hydroxide-oxide system

• The half-cell reactions are:

Cd (s) +  2OH^-  (aq) → Cd(OH)₂ (s)  +  2e^-                            E^ꝋ = -0.82 V 

NiO(OH) (s) + H₂O (l) + e^- → Ni(OH)₂ (s) +   OH^-  (aq)         E^ꝋ =  +0.38 V

• The cell generates an electromotive force of:

  • E_cell^ꝋ = E_reduction^ꝋ - E_oxidation^ꝋ

  • E_cell^ꝋ = (+0.38) - (-0.82) = +1.20 V

• The overall reaction is:

2NiO(OH) (s) + 2H₂O (l) + Cd (s) → 2Ni(OH)₂ (s) + Cd(OH)₂ (s)

• Cadmium is a toxic metal so the disposal of old NiCad cells is also an environmental issue

Lithium Cells

• Lithium ion cells power the laptop or mobile device you are probably reading this on

• The Noble Prize for Chemistry in 2019 was awarded to John B. Goodenough, M. Stanley Whittingham and Akira Yoshino for their work on lithium ion cells that have revolutionised portable electronics

• Lithium is used because it has a very low density and relatively high electrode potential

• Lithium cells consists of:

  • A positive lithium cobalt oxide electrode

  • A negative carbon electrode

  • A porous polymer membrane electrolyte

• The polymer electrolyte cannot leak since it is not a liquid or paste, which presents advantages over other types of cells

• The cell consists of a sandwich of different layers of lithium cobalt oxide and carbon

• When the cell is charged and discharged the lithium ions flow between the negative and the positive through the solid electrolyte

• The half-cell reactions on discharge are:

Li (s) →   Li⁺ (s)  +  e^–                                                                E^ꝋ = -3.05 V 

Li⁺ (s)  + CoO₂ (s)  +  e^– →   Li ⁺ (CoO₂) ^– (s)                E^ꝋ = +1 V

• The cell generates an emf of between 3.5 V and 4.0 V and the overall reaction is

         Li (s)  + CoO₂ (s)  →   Li ⁺ (CoO₂) ^– (s)                         E^ꝋ_cell ~ +3.5

• NiCad cells have a problem called the memory effect in which they gradually begin to lose their charge after repeated charge cycles when the cell is not fully discharged

  • The cells appear to 'remember' their lower state of charge

• Lithium-ion cells do not have this problem so can be topped up without any loss of charge

• Some of the problems with lithium ion cells:

  • A global shortage of lithium is likely to make lithium ion cells unsustainable as the current demand for lithium exceeds the supply

  • If cells are not recycled but thrown away in landfills, then a huge amount of lithium becomes lost to future generations

  • 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

Fuel Cells

• A fuel cell is an electrochemical cell in which a fuel donates electrons at one electrode and oxygen gains electrons at the other electrode

• These cells are becoming more common in the automotive industry to replace petrol or diesel engines

• As the fuel enters the cell it becomes oxidised which sets up a potential difference or voltage within the cell

• Different electrolytes and fuels can be used to set up different types of fuel cells

• An important cell is the hydrogen-oxygen fuel cell which combines both elements to release energy and water

• The fuel cell consists of:

  • A reaction chamber with separate inlets for hydrogen and oxygen gas

  • An outlet for the product - water

  • An electrolyte of aqueous sodium hydroxide

  • A semi-permeable membrane that separates the hydrogen and oxygen gases

• The half equations are

2H₂ (g) + 4OH^– (aq)  →  4H₂O (l) +  4e^–                     E^ꝋ = -0.83 V 

O₂ (g) +  2H₂O  +  4e^– →  4OH^– (aq)                      E^ꝋ = +0.40 V 

• The cell generates an electromotive force of:

  • E_cell^ꝋ = E_reduction^ꝋ - E_oxidation^ꝋ

  • E_cell^ꝋ = (+0.40) - (-0.83) = +1.23 V

• The overall reaction is found by combining the two half equations and cancelling the common terms:

 2H₂ (g) + O₂ (g)  →   2H₂O (l)

Benefits

• Water is the only reaction product, so fuel cells present obvious environmental advantages over other types of cells

• The reaction is the same as hydrogen combusting in oxygen, but since the reaction takes place at room temperature without combustion, all the bond energy is converted into electrical energy instead of heat and light

• There are no harmful oxides of nitrogen produced, which are usually formed in high temperature combustion reactions where air is present

• Fuel cells have been used on space craft, where the product can be used as drinking water for astronauts

Risks and problems

• Hydrogen is a highly flammable gas and the production and storage of hydrogen carries safety hazards

• 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

• Until a cheap way is found to make hydrogen, its widespread use in fuel cells will be limited

• Hydrogen has high energy density, that is, the amount of energy contained in 1g of the fuel is high compared to other fuels, but because it is a gas, its energy density per unit volume is low which means larger containers are needed compared to liquid fuels