The chemistry that underlies the depletion of ozone in the stratosphere by chlorofluorocarbon molecules is really interesting – an application of radical reactions in the atmosphere, initiated by high frequency UV-C and UV-B radiation.
The ozone layer and an ozone hole
The ozone layer refers to a region of the stratosphere approximately 15-35 km above the Earth’s surface with a higher concentration of ozone, O3. Ozone is formed when high energy UV-C radiation, with a wavelength of 100-280nm, is absorbed by diatomic oxygen molecules, O2, which photodissociate forming oxygen atoms. Almost all the very harmful UV-C radiation from the sun is absorbed by dioxygen in the stratosphere.
O2 ⇾ O + O
Oxygen atoms are radicals. Collision with a dioxygen molecule results in a molecule of ozone. Collision with ozone results in dioxygen.
O + O2 ⇾ O3
O + O3 ⇾ 2 O2
Ozone absorbs UV radiation in the wavelength range 200-310nm which accounts for a little UV-C and most UV-B radiation. UV-A radiation is transmitted to the Earth’s surface and is responsible for some skin cancers, sunburn and premature ageing.
O3 ⇾ O2 + O
Naturally, the rate of the reactions that create ozone are balanced by the rate of the reactions which destroy it.
However there are a number of man-made substances that play a significant role in removing ozone from the stratosphere – most notably chlorofluorocarbons, CFCs.
The depletion in the concentration of ozone in the ozone layer, particularly over the Artic and Antarctic, is of huge environmental concern.
CFCs such as CCl2F2 are very stable molecules chemically speaking as the C-Cl and C-F bonds are strong. They are not broken down in the troposphere or removed by rain.
CCl2F2 ⇾ CClF2 + Cl (C-Cl bond is weaker than C-F so is broken more easily)
The chlorine atoms / radicals formed react with ozone at a rapid rate in a cycle which regenerates the chlorine – it is essentially acting as a homogenous catalyst.
Timeline
1930: Thomas Midgely demonstrates the properties of CCl2F2 as a potential replacement for ammonia as a refrigerant – it was nontoxic, non-flammable, chemically stable, and with the low boiling and freezing points needed for a refrigerant. By 1970, 1000000 tonnes of CFCs had been manufactured worldwide for a variety of applications including aerosol propellants, blowing agents and dry cleaning solvents.
1974: Sherry Rowland publishes a paper predicting that CFCs would react with ozone in the stratosphere in the presence of uv radiation. By his calculation, each chlorine atom could destroy 100,000 molecules of ozone.
1984: The British Antarctic Survey confirms that there is a ‘hole’ in the ozone layer, based on the spectroscopic determination of the concentration of ozone in the stratosphere over several years. NASA had also been making similar measurements of ozone concentration by satellites but had failed to spot the alarming deterioration in ozone concentrations (the satellites collect so much data, more than could be processed, so computers analysing the data had been programmed to automatically ignore ‘impossibly inaccurate’ ozone measurements).
Why does the ozone hole develop over the Antarctic?
During the Antarctic winter, temperatures drop to below -80°C causing polar stratospheric clouds to form, providing surfaces onto which the chlorine reservoir molecules adsorb. Reactions between these reservoir molecules release Cl2 molecules which remain trapped in a vortex of air that circulates the Antarctic in winter.
ClONO2 + HCl → Cl2 + HNO3
In the Antarctic spring, the vortex starts to break up and the addition of sunlight (uv radiation) begins the catalytic cycle that destroys ozone.
Cl2 + uv radiation → Cl + Cl
The more variable temperatures in the Arctic mean that fewer and less stable stratospheric clouds are formed leading to less severe ozone depletion at this pole.
Why isn’t all the ozone destroyed by chorine atoms?
Some molecules known as chlorine reservoirs are able to effectively ‘store’ chlorine in the stratosphere, hence removing it from the ozone destroying cycle.
- HCl – formed when chlorine atoms react with methane
CH4 + Cl → CH3 + HCl
- ClONO2 – chlorine nitrate
NO2 + ClO → ClONO2
Both chlorine reservoir molecules are soluble in water and if carried into the troposphere they would be removed by rain.
CFC replacements
Hydrochlorofluorocarbons (HCFC) and hydrofluorocarbons (HFC) are widely used as CFC replacements. The molecules are (mostly) broken down in the troposphere so have a much lower / no ozone depleting potential, but they are known to contribute to global warming.
However, in recent times these CFC replacements have come under increasing environmental scrutiny themselves …
Calculating the frequency and wavelength of radiation needed to break a specific bond
This is a common exam question!
The average bond enthalpy of a C-F bond is +485 kJ mol-1. Calculate the wavelength of radiation needed to break this bond in nanometres.
- Calculate the bond enthalpy of a single bond
- Calculate the frequency of radiation needed to break this bond
- Calculate the wavelength of this frequency of radiation, convert to nm
Practice questions
1. State 2 effects of ultraviolet radiation on a molecule.
2. Calculate the bond enthalpy of a C-Br bond in kJ mol-1, given that the frequency of radiation needed to break the bond is 7.14 x 1014 Hz.
3. Calculate the concentration of ozone in the stratosphere in parts per million (ppm), given that its percentage by volume is 2.1 x 10-5 %.
4. Bromine atoms are able to catalyse the breakdown of ozone in the stratosphere. Use the equations shown below to explain how.
Br + O3 ⇾ BrO + O2
BrO + O ⇾ Br + O2
5. CCl2F2 photodissociates in the presence of high energy ultraviolet radiation to form two radicals. Write an equation using displayed formulae to represent this, showing the movement of electrons using curly arrows.
6. Calculate the wavelength of radiation (in nanometers) needed to break a single C-Cl bond, given that the average bond enthalpy for a C-Cl bond is +346 kJ mol-1.
7. Ozone is broken down in the stratosphere with both oxygen atoms and chlorine atoms able to catalyse the reaction. Describe fully the role these two species play in the process, including equations, and explain why chlorine atoms are the far greater threat.
8. Chloromethane is produced by naturally occurring forest fires. Suggest why chlorine radicals are not produced from chloromethane molecules in the troposphere.
9. Oxygen atoms can react with water in the troposphere to produce hydroxyl radicals, OH.
(a) write a balanced equation for this reaction
(b) draw a dot-cross diagram for a hydroxyl radical
(c) what does the term ‘radical’ indicate about the electronic structure of OH?
10. (a) Calculate the minimum frequency of radiation needed to break a C-Br bond, given that the average bond enthalpy of a C-Br bond is +290 kJ mol-1.
(b) How would the minimum frequency of radiation needed to break a C-Cl bond compare with that needed to break a C-Br bond? Explain your answer.
11. The overall equation for the formation of ozone is shown below.
3 O2(g) ⇾ 2 O3(g)
(a) Write two balanced equations to show how oxygen is converted into ozone in the stratosphere
(b) Explain why these reactions do not usually occur in the troposphere
12. 1-bromopropane is a possible substitute for chlorofluorocarbons used as solvents. Explain how manmade CFCs released into the troposphere damage the ozone layer and suggest why 1-bromopropane is likely to be less of a problem.
13. (a) Give one property of CFCs that makes them suitable for use as refrigerant gases.
(b) Suggest two factors that need to be considered in designing a suitable replacement for a CFC.
Answers
1. UV radiation increases the electronic energy of a molecule and can cause homolytic bond fission of the C-Cl bond.
10. (b) C-Cl bonds are stronger than C-Br bonds (Cl is a smaller atom so there is a greater attraction to the shared pair of electrons and the bond is shorter), so a greater frequency of radiation would be needed.
11. (a) O2 ⇾ O + O and O + O2 ⇾ O3
(b) The high frequency UV-B radiation which is needed to break the bond in the dioxygen molecule is only present in the stratosphere.
12. CFCs released into the troposphere are not broken down as the frequency of radiation is not high enough to cause the strong C-Cl and C-F bond to break, so they make their way to the stratosphere. The high frequency UV-B radiation in the stratosphere causes the CFCs to photodissociate, forming Cl radicals which are able to catalyse the breakdown of ozone into dioxygen. Since the C-Br bond is weaker than the C-Cl bond, 1-bromopropane is more likely to be broken down in the troposphere, and hence less of a concern with regards to ozone depletion.
13. (a) volatile / unreactive / nontoxic
(b) volatility / flammability / toxicity / cost of manufacture / ease of disposal