Ozone Revisited (DP IB Chemistry: HL)

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

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 117^o

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

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 on the oxygens in ozone

• 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

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) and you can show the second step where it reacts with another molecule of ozone, producing oxygen and regenerating the NO⋅ radical:

NO₂⋅ (g) + O₃ (g) → NO⋅ (g) + 2O₂ (g)

• An alternative to the second step shows the NO₂⋅ reacting with an oxygen radical to produce the same products but in a different stoichiometry

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