Back
to top
Ozone formation
Watch this threadPage 1 of 1
Go to first unread
Skip to page:
Rexx18
Badges:
10
Rep:
?
You'll earn badges for being active around the site. Rep gems come when your posts are rated by other community members.
#1
Let * represent the free-radical.
O2 + hv --> O* + O*
O2 + O* --> O3
How come there's a radical on the LHS but not one on the RHS? Is it something to do with the structure of ozone?
O2 + hv --> O* + O*
O2 + O* --> O3
How come there's a radical on the LHS but not one on the RHS? Is it something to do with the structure of ozone?
0
reply
MexicanKeith
Badges:
12
Rep:
?
You'll earn badges for being active around the site. Rep gems come when your posts are rated by other community members.
#2
Report
#2
(Original post by Rexx18)
Let * represent the free-radical.
O2 + hv --> O* + O*
O2 + O* --> O3
How come there's a radical on the LHS but not one on the RHS? Is it something to do with the structure of ozone?
Let * represent the free-radical.
O2 + hv --> O* + O*
O2 + O* --> O3
How come there's a radical on the LHS but not one on the RHS? Is it something to do with the structure of ozone?
When it recombines with O2 to form ozone this extra electronic energy can be converted into extra vibrational energy in the ozone.
The ozone can then lose the vibrational energy by colliding with another molecule (eg N2 in the atmosphere, this is often given the symbol "M" which just refers to a generic third molecule)
so the step could be written in the following ways
O2 + O* ---> O3
OR
O2 + O* ----> O3*
O3* + M ---> O3 + M
OR
O2 + O* + M ---> O3 + M
so the reason a radical isn't present on the right hand side is simply because the ozone can lose the extra energy by colliding with a third molecule!
1
reply
X
Page 1 of 1
Go to first unread
Skip to page:
Quick Reply
