IB Chemistry HL

Revision Notes

14.2.2 Ozone Revisited

Ozone Revisited

  • We have seen previously that ozone is a molecule with two resonance structures leading to a resonance hybrid

Lewis structures for ozone, downloadable IB Chemistry revision notes

The two Lewis resonance structures for ozone

  • The central oxygen atom has three electron domains and a lone pair, so the domain geometry is triangular planar and the molecular geometry is bent linear
  • The presence of the lone pair repels the bonding pairs more strongly so the bond angle is reduced to 117o

Molecular-structure-of-ozone, downloadable IB Chemistry revision notes

The molecular structure of ozone

  • The bond order for each bond in ozone is

bond order in O3 = total number of O3 bonding pairs ÷ total number of positions = 3 ÷ 2 = 1.5

  • This gives a polar molecule with bonds that are weaker than the double bond in oxygen molecules

Oxygen and ozone, downloadable IB Chemistry revision notes

The structure of oxygen and ozone

  • You would expect O-O bonds to be non-polar as the atoms have the same electronegativity; this is correct, but overall the molecule is polar due to the uneven distribution of electron cloud charge
  • The formal charge on the Lewis structures show that the electrons are unevenly distributed

FC= (number of valence electrons) – ½(number of bonding electrons) – (number of non-bonding electrons)

FC (oxygen A) = (6) – ½(2) – (6) = -1

FC (oxygen B) = (6) – ½(6) – (2) = +1

FC (oxygen C) = (6) – ½(4) – (4) = 0

Formal charges of the oxygen in ozone, downloadable IB Chemistry revision notes

Formal charges on the oxygens in ozone

Catalytic Depletion

  • The bonding and structure of ozone is key to understanding how the catalytic depletion of ozone occurs in the stratosphere
  •  High energy UV radiation in the stratosphere breaks the oxygen-oxygen double bond creating oxygen atoms

O2 (g) → O⋅ (g) + O⋅ (g) ∆H +ve, UV light, λ < 242 nm

  • These oxygen atoms have unpaired electrons- they are known as free radicals
  • The free radicals are highly reactive and quickly attack oxygen molecules forming ozone in an exothermic reaction, which raises the temperature of the stratosphere

OZONE FORMATION         O⋅ (g) + O2 (g) → O3 (g) ∆H – ve

  • Ozone requires less energy to break than oxygen
  • It produces an oxygen molecule and an oxygen free radical:

OZONE DEPLETION          O3 (g) → O⋅ (g) + O2 (g) ∆H +ve, UV light, λ< 330 nm

  • The radical reacts with another ozone molecule making two molecules of oxygen in an exothermic reaction

OZONE DEPLETION         O3 (g) + O⋅ (g) → 2O2 (g) ∆H – ve

  • The temperature in the stratosphere is maintained by the balance of ozone formation and ozone depletion in a process known as the Chapman Cycle
  • It is not a closed system as matter and energy flow in and out, but it is what is called a steady state

The Chapman cycle, downloadable IB Chemistry revision notes

The Chapman cycle

Catalytic Depletion

  • The two main man made culprits that accelerate the depletion of ozone are nitrogen oxides and CFCs
  • Nitrogen monoxide, NO, is produced from the high temperatures inside internal combustion engines
  • If you count the valence electrons in nitrogen monoxide (5 + 6 =11), the odd number tells you it is a free radical as it has an unpaired electron
  • The nitrogen monoxide reacts with ozone forming oxygen and a nitrogen dioxide radical

NO⋅ (g) + O3 (g) → NO2⋅ (g) + O2 (g)

  • The nitrogen dioxide produced is also a free radical (it has 5 + 6 + 6= 17 electrons)

NO2⋅ (g) + O⋅ (g) → NO⋅ (g) + O2 (g)

  • The nitrogen monoxide is regenerated so it has a catalytic role in the process
  • Combining the two equations and cancelling out the NO⋅ and NO2⋅ and you arrive at the overall depletion of ozone

O3 (g) + O⋅ (g) → 2O2 (g)

  • A similar process happens with CFCs
  • The C-Cl bond in the CFCs is weaker than the C-F bond and breaks more easily in the presence of UV light creating chlorine radicals

CCl2F2 (g) + UV → CClF2⋅ (g) + Cl⋅ (g)

  • The chlorine radicals attack ozone and are regenerated at the end of the cycle

Cl⋅ (g) + O3 (g) → ClO⋅ (g) + O2 (g)

ClO⋅ (g) + O⋅ (g) → Cl⋅ (g) + O2 (g)

  • Once again a molecule of ozone has been destroyed by a catalytic free radical
  • The net effect of these reactions is that these pollutants have created an imbalance in the natural ozone cycle leading to an overall depletion in stratospheric ozone
  • CFCs are greatly damaging to stratospheric ozone and have been largely replaced by safer alternatives following the 1985 Montreal Protocol
  • The depletion of ozone has allowed greater amounts of harmful UV light to reach the surface of the Earth
  • UV light has been linked to greater incidence of skin cancer and cataracts as well as the destruction of phytoplankton and reduced plant growth

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