Equations + industry

• Flue gas removal 


  CaO(s) + 2H20(l) + SO2(g) -> CaSO3(aq) + 2H20(l)

  CaCO3(s) + 2H20(l) + SO2 -> CaSO3(aq) + 2H20(l) + CO2(aq)
• Cl- & F- and sulphuric acid


  Cl- and F- are too weak reducing agents to cannot reduce sulphur and the reaction only produces the steamy fumes of the hydrogen hallide:

  l NaCl + H2SO4 -> HCl + NaHSO4

  Observations:

  Steamy fumes (that turn litmus paper red), white solid (sodium bisulfate)   
• Bromide ions and sulphuric acid 


  l Bromide only strong enough to reduce sulphur from +6 to +4

  2Br-    Br2  +  2e-

  H2SO4 + 2H+ + 2BR- _> Br2 + SO2 + 2H20

  Combines to make:

  l H2SO4 + 2H+ + 2Br- -> Br2 + SO2 + 2H20

  Observations:

  Orange fumes (bromine), (colourless odourless SO2 has also been produced)
• Iodide ions and sulphuric acid

  Iodide can reduce sulphur from +6 to +4, 0, -2

  Sulphur dioxide, sulphur, hydrogen sulphide    

  l H2SO4 + 8H+ + 8I- -> 4I2 + H2S + 4H20

   

  Observations:

  Purple fumes (iodine vapour), dark grey solid (iodine which condenses at the top of the tube), choking gas with bad egg smell produced (H2S)
• Uses of chlorine/chlorate

   Reacts with water:

  Cl2(g) + H20(l) <-> HCl(aq) + HCLO(aq)

  HClO decomposes in the presence of sunlight:

  2HClO(aq) -> 2HCl(aq) + O2(g) (sunlight)

   

  OVERALL:

  2H20(l) + 2Cl2(g) <-> 4HCl(aq) + O2(g)

   

  l In presence of UV light:

  2Cl2 + 2H20 -> 4HCl + O2 
• l BLEACH PRODUCTION:

  Cold aqueous sodium hydroxide

  2NaOH + Cl2 -> NaClO + NaCl + H20
• Lithium cobalt oxide batteries

  • LiCoO2 batteries were the first lithium ion batteries invented.
  • They have one electrode made of LiCoO2 and one made of graphite.
  • The two half equations are:
  • Li+ + e- ⇌ Li(s), Eθ = −3.04V
  • Li+ + CoO2 + e- ⇌ LiCoO2, Eθ = 0.56V
  • The Li+/Li(s) cell has a more negative Eθ so is reversed.
  • The overall cell potential is: Eθcell = 0.56 − (−3.04) = 3.6V.
  • To reverse the reaction, you apply an external current of at least 3.6V to push the reactions in the opposite directions.
• Alkaline fuel cell

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

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

  overall: 2H2 (g) + O2 (g)  →   2H2O (l)           Eꝋ = +1.23 V
• Acid fuel cell

  Negative electrode:   H2 (g) →  2H+ (aq) +  2e–   Eθ = 0.00 V

  Positive electrode:   O2 (g) +  4H+ (aq) +  4e– →  2H2O (l)   Eθ = +1.23 V 

  Overall:

  2H2 (g) + O2 (g) + 4H+ (aq) +  4e– →   2H2O (l) +  4H+ (aq) + 4e– 

   2H2 (g) + O2 (g)  →   2H2O (l)           Eθ = +1.23 V
• Advantages of fuel cells

  Advantages of fuel cells

  • Fuel cells have a high efficiency.
  • This means that they get more energy out of the same amount of fuel than less efficient devices.
  • Their only by-product is water (no CO2(g) is produced).
  • They don’t need to be recharged - they will work as long as the fuel is supplied
• Disadvantages of fuel cells

  Disadvantages of fuel cells

  • Hydrogen is flammable and so must be stored very carefully.
  • Energy is needed to produce the reactants of hydrogen and oxygen.
  • This energy usually comes from fossil fuels, so there are some CO2(g) emissions.
• Vanadium catalyst- heterogenous 


  V2O5 + SO2 --> V2O4 + SO3

  V2O4 + 1/2O2 --> V2O5