Terms in this set (30)
Molecules in the atmosphere are continually moving and colliding with
Chemical reactions in atmosphere dependent on these collisions
Molecular collisions in atmosphere determined by temperature and concentration—the higher the concentration, the more collisions
—the higher the temperature, the more collisions
Atmosphere continually illuminated during daylight
Chemical reactions which occur due to the absorption of photons
determining the composition of the atmosphere itself
contribute to air pollution
Nitrogen dioxide, NO2
one of the most photochemically active species in the atmosphere
contains an unpaired electron
electronically excited molecule
a molecule which absorbs a photon of light with energy hυ,
remit electron or form products
∙NO2(g) → ∙NO(g) + ∙O∙(g)
another mechanism of formation of radicals
a molecule absorbs an ultraviolet photon and produce s two free radicals
O2(g) → ∙O∙(g) + ∙O∙(g)
radicals are highly reactive and short-lived
a pollutant in the troposphere
resonance structures, bond angle 116.8o
Ozone less stable than oxygen—proof is the Enthalpy of Formation
3 O2(g) → 2 O3(g) ΔHof of 142.3 kJ/mol
Comparison to O2 hυ
∙O∙(g) + ∙O∙(g) → O2(g) ΔHof of 0 kJ/mol
The heat of formation of an element in its standard state is 0.0 kcal/mol
The positive enthalpy means that the formation of ozone requires energy.
stratospheric ozone layer
altitude of 25 to 30 km, O3 concentration 10 ppm
ozone concentration maintained by a reaction that destroys ozone at the
same rate at which it is produced.
Natural Destruction of O3
O3(g) + ∙O∙(g) → 2 O2(g)
O3 destruction reaction slow
ozone molecules and oxygen atom radicals are very low concentrations
In order to react the O3 and O∙ have to collide
If the concentrations are low of both species, then they do not often randomly collide with one another, so it takes the reaction a long time
because it takes a long time for the reactants to collide with one another
Even though slow, ozone breaks down fast enough to maintain a balance with the ozone-forming reaction, so that under normal circumstances the concentration of ozone in the stratosphere remains constant.
Ozone destroyed & formed each day
3 X 10 to the 8 tons of stratospheric ozone
stratospheric ozone prevents damaging ultraviolet radiation from
reaching the earth's surface
O3 decomposition reaction
absorbs wavelengths in the range of 200 to 310 nm
O3(g) → ∙O∙(g) + O2(g)
Stratospheric ozone prevents 95% to 99% of the sun's ultraviolet radiation from reaching the earth's surface
Skin cancer caused
Photons 200 to 310-nm range
Stratospheric ozone depletion
serious threat to Earth-- every 1% decrease in the stratospheric ozone is
an additional 2% damaging radiation
small molecules: halogen atoms bonded to the central carbon atoms
Properties of CFCs
nonflammable, relatively inert, volatile, readily liquefied, and nontoxic.
used as coolants for refrigeration, in foam plastics manufacture, as aerosol propellants, and as industrial solvents.
same properties that make CFCs useful for a wide arrange of consumer products and industrial uses accounts for their destructive effect on the stratospheric ozone
Gaseous CFCs released into the atmosphere persist for a very long time in troposphere due to their chemical nonreactivity. A CFC molecule survives roughly 100 years in the troposphere.
CFCs eventually mix into stratosphere, where they decompose by
high-energy solar radiation.
Photodissociation of CFCs
decomposition products of CFCs lower stratospheric ozone concentration
CF2Cl2(g) → CF2Cl ∙(g) + Cl ∙ (g)
chain reaction mechanism of O3 destruction involving Cl ∙ radical from CFC
Step 1: Cl ∙(g) + O3(g) → ClO∙(g) + O2(g)
Thus, an ozone molecule has been destroyed. If this were the only reaction that particular CFC molecule caused, there would be little danger to the ozone layer. However, once Step 1 has occurred twice,
the two ClO∙ radicals produced can react further.
Step 2: ClO∙+ ClO∙ → ClOOCl
Step 3: hυ
ClOOCl(g) → ∙ClOO + ∙Cl
Step 4: hυ
∙ClOO → ∙Cl + O2(g)
Net reactions: 2 O3(g) → 3 O2(g)
chain reaction increases rate of stratospheric ozone destruction,
but it does not affect the rate of stratospheric ozone formation
∙Cl catalyst speeds up ozone destruction reaction
is not used up by the reaction
lowers the energy requirements for ozone destruction, O3 does not absorb the UVC light
CONTINUES to be around to do the same for other reactions
a single chlorine atom destroys 100,000 molecules of O3 before it is inactivated or returned to the troposphere (probably as HCl).
First lab evidence that CFCs could deplete ozone layer
In 1974 F. Sherwood Rowland and Mario Molina of UC Irvine show CFCs capable of depleting the ozone layer in the atmosphere.
first confirmation of stratospheric ozone depletion and increased ultraviolet intensity
September 10, 1999
Antarctic ozone hole
Discovered 1985 by British team. 1986, NASA chooses 30 year old Susan Solomon's team to find proof of what is causing ozone hole. Solomon's team goes to Antarctica and shows high concentrations of ClOOCl, to support theory of chlorine radical chain reactions atoms in Winter moon beam light over Antarctica
Photodissociation of Cl2
In the dark Antarctic winter, a vortex of intensely cold air containing ice
crystals, unique to the region, builds up.
On the surfaces of these crystals additional reactions produce hydrogen chloride and chlorine nitrate (ClONO2).
They can react with each other to form chlorine molecules.
HCl (g) + ClONO2 (g) → Cl2 (g) + HNO3 (g)
When sunlight returns in the spring, the Cl2 molecules are readily
photo dissociated into chlorine atoms, which can then become involved in the ozone destruction reactions.
Cl2 (g) → 2 Cl∙ (g)
Dobson units, the UV index
1 DU = 2.69 X 1010 molecules/cm3
Montreal Protocol on Substances that Deplete the Ozone Layer 1987
Protocol has been amended four times—London (1990), Copenhagen (1992), Montreal (19997), and Beijing (1999)—and has now been signed by more than 150 countries.
It is part of Title VI of the Clean Air Act in the United States
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