5.4.4 Commercial Cells

• Electrochemical cells can be used as a commercial source of electrical energy
• Cells can be non-rechargeable (irreversible), rechargeable or fuel cells
• Type of cells used in commercial applications depend on
  • the voltage required
  • the current needed
  • the size of the cell
  • the 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 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

A Daniell cell

• 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 emf of 1.1 V and the overall reaction is

Zn (s)  + CuSO₄ (aq)→ Cu (s)  + ZnSO₄ (aq)      E^ꝋ_cell = +1.10 V

• 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, consisting of
  • a zinc casing which acts as the negative electrode
  • 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-carbon cell

• 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 emf of 1.50 V and the overall reaction is

2NH₄⁺ (aq) + Zn (s)  → 2NH₃ (g) + H₂ (g)  + Zn²⁺ (aq)    E^ꝋ_cell = +1.50 V

• As the cell discharges, the zinc casing eventually wears away and the corrosive contents of the electrolyte paste can leak out, which is an obvious disadvantage of zinc-carbon cells
• The cell provides a small current and is relative cheap compared to other cells
• Extra long life cells have a similar chemistry, but supply a higher current and use zinc chloride in the paste; they are suitable for torches, radios and clocks
• Another variation on the cell uses an alkaline paste in the electrolyte and they have a much longer operating life, but are noticeably more expensive than regular zinc-carbon cells

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 and lithium cells which are 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

A lead-acid battery

• 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

• The cell generates an emf of about 2 V and the overall reaction is

PbO₂ (s) +  4H⁺ (aq) +  2SO₄^2- (aq) +  Pb (s) →  2PbSO₄ (s)  + 2H₂O (l)       E^ꝋ_cell = +2.06 V

• 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 materials: 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 overall reaction in the cell is

2NiO(OH) (s) + 2H₂O (l) + Cd (s) → 2Ni(OH)₂ (s) + Cd(OH)₂ (s)         E^ꝋ = +1.2 V

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