| Pages:
1
2 |
morganbw
National Hazard
Posts: 561
Registered: 23-11-2014
Member Is Offline
Mood: No Mood
|
|
Quote: Originally posted by S.C. Wack  |
Or not.
How do you explain chloroform being missing from Montreal and ODP lists? Stupidity? Ignorance? Bribes from CHCl3 users and manufacturers? Being
naturally occurring, not photoreactive, and with a short atmospheric lifetime and general inability to reach the ozone layer?
Oxidation by ozone would actually be a great thing, since we'd be talking about ground level ozone aka pollution. Maybe China should be stocking up
for next time they hold some international prestige event so they don't have to shut everything off? |
I think he is going by memory and is probably confusing it with methyl chloroform. Good chance anyway.
|
|
|
S.C. Wack
bibliomaster
Posts: 2419
Registered: 7-5-2004
Location: Cornworld, Central USA
Member Is Offline
Mood: Enhanced
|
|
Yes, I'm just being a dick with pent-up frustration at all the threads and posts from many of our um...newer...and/or more prolific members...
The majority of the chloroform in the atmosphere is natural and oceanic in origin, which is a good enough excuse for me to...well actually I'd never dispose of chloroform...on purpose.
http://www.tellusb.net/index.php/tellusb/article/view/14614
[Edited on 14-12-2014 by S.C. Wack]
|
|
|
AJKOER
Radically Dubious
Posts: 3026
Registered: 7-5-2011
Member Is Offline
Mood: No Mood
|
|
One reference (see http://scorecard.goodguide.com/chemical-profiles/def/odp.htm...):
"Ozone depleting substances (ODS), including chlorofluorocarbons (CFCs), halons, and several other chemicals, are responsible for thinning the
stratospheric ozone layer. When these substances reach the stratosphere, UV radiation from the sun breaks them apart to release chlorine or bromine
atoms which react with ozone, starting chemical cycles of ozone destruction that deplete the ozone layer. One chlorine atom can break apart more than
100,000 ozone molecules, while a bromine atom can destroy about 4,000,000 ozone molecules"
Then this reference on page 13.4 (https://www.google.com/url?sa=t&source=web&rct=j&... ):
"As such, the total lifetime must take into account all of the processes determining the removal of a gas from the atmosphere, including photochemical
losses within the troposphere and strato sphere (typically due to photodissociation or reaction with OH), heterogeneous removal processes, and
permanent removal following uptake by the land or ocean."
Followed by the authors comment on page 13.5:
"From this total atmospheric lifetime, together with the evaluated loss lifetimes of CH3CC13 due to the ocean (about 85 years, with an uncertainty
range from 50 years to infinity; see Butler et al., 1991) and stratospheric processes (40 ± 10 years), a tropospheric lifetime for reaction with OH
of 6.6 years can be inferred (±25% ). The lifetimes of other key gases destroyed by OH (i.e., CH4, HCFCs, and hydrofluorocarbons [HFCs]) can then be
inferred relative to that of methyl chloroform (see, e.g., Prather and Spivakovsky, 1990) "
And conveniently, per Table 13.1 on page 13.6, CH3CCl3 and CHCl3 are listed consecutively, with the relative hazard of Chloroform specifically, only
one tenth of methyl chloroform. So apparently CH3CCl3 is more readily prone (possibly owing, in part, to its relative molecular size increasing the
likelihood of a collision, I would guess) under atmospheric conditions to attack by the hydroxyl radical primilary (from the action of intense uv
radiation on water vapor in the upper atomsphere), as was suggested by S.C. Wack reference above, than CHCl3 followed by photodissociation to the Cl
radical, a catalytic agent in ozone destruction.
[Edit] Here is some background comments per Wikipedia http://en.m.wikipedia.org/wiki/Ozone_depletion:
"The Cl and Br atoms can then destroy ozone molecules through a variety of catalytic cycles. In the simplest example of such a cycle,[4] a chlorine
atom reacts with an ozone molecule, taking an oxygen atom with it (forming ClO) and leaving a normal oxygen molecule. The chlorine monoxide (i.e., the
ClO) can react with a second molecule of ozone (i.e., O3) to yield another chlorine atom and two molecules of oxygen. The chemical shorthand for these
gas-phase reactions is:
Cl· + O3 → ClO + O2: The chlorine atom changes an ozone molecule to ordinary oxygen
ClO + O3 → Cl· + 2 O2: The ClO from the previous reaction destroys a second ozone molecule and recreates the original chlorine atom, which can
repeat the first reaction and continue to destroy ozone."
And finally, my comment of CHCl3 ceasing to destroying O3 occuring after some 50,000 molecules of ozone are consumed should be most likely be revised
down to say 10,000 (assuming the above cited 100,000 being applicable to CH3CCl3 as a source of the Cl radical, and the noted 1/10 relativity for
CHCl3 specifically) or fewer, which some may still find troubling.
[Edited on 15-12-2014 by AJKOER]
|
|
|
| Pages:
1
2 |
|
![[*]](images/xpblue/default_icon.gif){kind=icon}
