From: sbharris@ix.netcom.com (Steven B. Harris )
Subject: Re: Physiological NO
Date: 09 Sep 1995
Newsgroups: sci.life-extension

In <pproctor.405.0069A78F@sam.neosoft.com> pproctor@sam.neosoft.com
(Peter H. Proctor) writes:

>2) There is some evidence that NO is bound in biological systems.
>Typically, this is in the form of a complex with a sulfhydryl compound,
>e.g. cysteine. Interestingly, the Procter and Gamble corp just got a US
>Patent to use acetyl cysteine to treat wrinkles.


ROFL!  The world's most useless patent.  I know a cosmetics person who
looked at that years ago and rightly concluded that no woman would use
a face cream that smelled like rotten eggs, no matter how good it made
her look after months of use!

>     Nitric oxide may also be present in biological systems as a
>hydroxylamine.  The half-life of such complexes is significantly longer
>than NO itself.  Even though NO probably does mean NO, some folks like me
>still hedge a bit and call it "endothelium-derived relaxing factor" or
>EDRF.
>
>Peter H. Proctor, PhD, MD


   Well not the ones I know.  I think it's pretty well accepted that NO
is EDRF, period.  If NO then gets metabolized to other things that also
cause relaxation, that's okay.  Growth hormone didn't stop being growth
hormone when we discovered somatomedin C/ IGF-1.

                                          Steve Harris

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.life-extension,sci.med
Subject: Re: Free radicals
Date: 27 Jul 1998 06:51:28 GMT

In <35bb562f.5802923@news.mindspring.com> Stuart writes:
>			  The electro-chemical metabolic, immune and
>detoxification reactions of the body are not presently structured into
>an organized body of knowledge. However we know increased
>levels of active oxidants, or free radicals, create oxidative stress,
>resulting in metabolic dysfunction. Genes of antioxidant enzymes
>have been cloned and characterized. Laboratories study these genes
>to engineer organisms for increased tolerance to oxidative stress.
>Thus the question - are atoms within the body at their lowest possible
>state of electro-excitation when they enter into metabolic reactions -
>or are these (eg. O-16) atoms truly identical?



Comment:

   The nuclei of such atoms are identical, which is what I hope I said,
but maybe didn't.

   Free radicals are not all bad.  They're your body's main signal
molecules for fighting infection, healing wounds, responding to damage.
(think about nitric oxide, for example).   Think of them like a old
time street cop's whistle.  You can't just damp them all out without
paying a price for it.

    Lots of experiments have been done with mammals to see if sopping
up more of their free radicals makes them age more slowly.  So far, all
are basically negative.  So it's not the greatest theory of aging.

   Free radicals ARE major players in inflammation, and sometimes the
body doesn't know how much inflammation is too much.  In the young,
inflammation is designed to protect you from infection after external
injury (fighting for that female, or to get that food).  In the old
animal or human, however, the damage may NOT be from chasing food or
women, but from age related wear and tear, and in that case the
inflammation is inappropriate (not that evolution cares, since it
didn't design your systems to work well in old age).  It's probably
here that antioxidents shine, and are good for you.  Also in injuries
that are huge (as to your entire brain, or massive body injuries) which
evolution didn't really design an appropriate antioxident response for,
since you weren't expected to survive anyway.  And antioxidents and
antiinflammatories are useful as adjunct treatments for infection in
moderation, if antibiotics take up some of the slack.  Evolution
doesn't know about antibiotics, either, and left to her own devices,
mother nature generally uses too much inflammation for the optimum job,
if antibiotics are also in there working.  Doctors found that out with
infant meningitis, where they found that steroids along with
antibiotics greatly improved outcome.  But there are many more
examples.


                                    Steve Harris, M.D.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.life-extension,misc.health.alternative
Subject: Re: anti-oxidants and cancer prevention
Date: 12 Feb 1999 09:04:51 GMT

In <79gm5d$osc$1@mozart.jlc.net> "Kate" <klindner@jlc.net> writes:

>  But is there an evolutionary reason to expect our system to crank up
>  the production of some antioxidants and shut off others? Unless there
>  is something special about these particular antioxidants? If the are
>  cheaper in evolutionary currency( available precursors, less
>  endothermic )? Or if they interfere less with the legitimate function
>  of free radicals in the immune system?


    I rather think the latter.  Evolution will have tended to optimize
you for a free-radical tone which is able to best deal with infections
and injuries (free radicals signal many inflammatory responses
necessary or healing) in an era where there are no antibiotics, and
your age is reproductive or younger.

   That might not be the best tone for a modern world where there are
fewer injuries, existant antibiotics, and older people who are
suffering more cumulative free radical effects.  Evolution is smart
over the eons, but it does't deal with rapid changes in environment
very well.  And since rapid change in environment is just another
aspect of human cultural development these days, we find ourselves
fighting the "wisdom" of evolution a lot, in medicine.

  Something the tree-huggers and natural hygeine folks have a hard time
understanding-- but then you don't see them running around naked and
eating raw things in the forest that much.  Generally I think they're
all too busy taking vitamin pills with more vitamin E than they could
get from a gallon of safflower oil, and sitting at their computers in
their centrally heated homes, putting out propaganda on the world wide
electronic web.  Or driving their petroleum-fueled cars to
demonstrations in which they shout into megaphones.

                                        Steve Harris, M.D.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.life-extension,misc.health.alternative
Subject: Re: anti-oxidants and cancer prevention
Date: 15 Feb 1999 06:22:51 GMT

In <7a85n7$9d3$1@mozart.jlc.net> "Kate" <klindner@jlc.net> writes:

>Steven B. Harris <sbharris@ix.netcom.com> wrote in message
>
>> I rather think the latter. Evolution will have tended to optimize you
>> for a free-radical tone which is able to best deal with infections and
>> injuries (free radicals signal many inflammatory responses necessary or
>> healing) in an era where there are no antibiotics, and your age is
>> reproductive or younger.
>
>
> This still doesn't explain how long-lived birds, that apparently
> produce few free radicals, manage to cope with infection.
>
>
>
     Kate



Answer:

   I don't know that they do, at least per young mitochondrion.  I
think rather that the rate of increased production that comes with
mitochondrial aging, is slowed in birds.  Or course, (if you'll forgive
me), it's a chicken and egg problem.

    If birds really do produce more free radicals per cell at rest (and
they certainly do if you're talking about raw numbers, since this is
tied in with numbers of mitochondria per cell, and specific metabolic
rate). then they and all small animals have no problems generating
enough for immune responses.  The question is how WE large and slow
metabolism mammals manage.  Obviously our immune systems are just more
sensitive to free radical signalling.  But we've evolved for the rate
of free radicals we make in various situations, and messing too much
with it, no matter what it is, is likely to be ungood, unless you have
a very good reason to think that either evolution is working against
you, or that you have something going for you that evolution doesn't
know about.  Probably you can do some damping, and compensate with
lower injury rates (safety stuff), antibiotics and other medical
support when needed, and (of course) with greater age, which gives you
more free radical production to begin with.  The trick is finding the
balance.

   BTW, this whole question is a subset of a very nasty problem that
the doctors of the next 20-30 years are going to have to wrestle with
every day.  Increasingly, we are getting control of the basic
mechanisms of inflammation, apoptosis, and immune activation--- all
very much tied in with free radical cascades as signals.  All these
cause tissue damage in injury and ischemia, and even in infection.  And
they may cause some of what we call aging.  But they also aid in injury
and infection control, and in healing.  So how much do you damp them
out at any one time?  Answer: as much as you can get away with.  But
that's a thing very hard to judge.  All you know is when you've gone
too far, and septic shock and multi-system organ failure are both
cascades that kill you like a house of cards.  It's going to be
increasingly a high wire act in ICUs, with literal moment by moment
monitoring and modulating of free radical tone of various kinds, and
tissue response to inflammation and infection control.  It will take a
lot of sensors and a lot of computers.  It's the kind of thing that is
going to be way beyond the skills of the fastest human mind.

                                       Steve Harris, M.D.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.life-extension,misc.health.alternative
Subject: Re: anti-oxidants and cancer prevention
Date: 17 Feb 1999 04:03:16 GMT

In <36c9afaf.19195140@nntp.uio.no> alf.christophersen@basalmed.uio.no
(Alf Christophersen) writes:

>>A search of avian mitochondria on pub med produced several abstracts
>>that showed avian mitochondria produce fewer free radicals than that of
>>mammals for similar energy production.


     Yep, looking it up I find that, by golly, you're right.  However,
it still could be a coincidence (a difference between orders that has
nothing to do with anything, like feathers vs hair).  I'll be much more
impressed when they find that bats, which have aging and metabolic
rates like birds, have the same mitochondrial efficiency, even though
they are mammals.

                                        Steve Harris

  Ann N Y Acad Sci 1998 Nov 20;854:224-38


Mitochondrial free radical production and aging in mammals and birds.

Barja G

Department of Animal Biology-II (Animal Physiology), Faculty of Biology,
Complutense University, Madrid, Spain.


[Medline record in process]

The mitochondrial rate of oxygen radical (ROS) production is negatively
correlated with maximum life span potential (MLSP) in mammals following
the rate of living theory. In order to know if this relationship is more
than circumstantial, homeothermic vertebrates with MLSP different from
that predicted by the body size and metabolic rate of the majority of
mammals (like birds and primates) must be studied. Birds are unique
because they combine a high rate of basal oxygen consumption with a high
MLSP. Heart, brain, and lung mitochondrial ROS production and free
radical leak (percent of total electron flow directed to ROS production)
are lower in three species of birds of different orders than in mammals
of similar body size and metabolic rate. This suggests that the capacity
to show a low rate of ROS production is a general characteristic of
birds. Using substrates and inhibitors specific for different segments of
the respiratory chain, the main ROS generator site (responsible for those
bird-mammalian differences) in state 4 has been localized at complexes I
and III in heart mitochondria and only at complex I in nonsynaptic brain
mitochondria. In state 3, complex I is the only generator in both
tissues. The results also suggest that the iron-sulphur centers are the
ROS generators of complex I. A general mechanism that allows pigeon
mitochondria to show a low rate of ROS production can be the capacity to
maintain a low degree of reduction of the ROS generator site. In heart
mitochondria, this is supplemented with a low rate of oxygen consumption
physiologically compensated with a comparatively higher heart size. A low
rate of free radical production near DNA, together with a high rate of
DNA repair, can be responsible for the slow rate of accumulation of DNA
damage and thus the slow aging rate of longevous animals.

PMID: 9928433, UI: 99127394

----------

J Bioenerg Biomembr 1998 Jun;30(3):235-43


Localization at complex I and mechanism of the higher free radical
production of brain nonsynaptic mitochondria in the short-lived rat than
in the longevous pigeon.

Barja G, Herrero A

Department of Animal Biology-II (Animal Physiology), Faculty of Biology,
Complutense University, Madrid, Spain.

Free radical production and leak of brain nonsynaptic mitochondria were
higher with pyruvate/malate than with succinate in rats and pigeons.
Rotenone, antimycin A, and myxothiazol maximally stimulated free radical
production with pyruvate/malate but not with succinate. Simultaneous
treatment with myxothiazol plus antimycin A did not decrease the
stimulated rate of free radical production brought about independently by
any of these two inhibitors with pyruvate/malate. Thenoyltrifluoroacetone
did not increase free radical production with succinate. No free radical
production was detected at Complex IV. Free radical production and leak
with pyruvate/malate were higher in the rat (maximum longevity 4 years)
than in the pigeon (maximum longevity 35 years). These differences
between species disappeared in the presence of rotenone. The results
localize the main free radical production site of nonsynaptic brain
mitochondria at Complex I. They also suggest that the low free radical
production of pigeon brain mitochondria is due to a low degree of
reduction of Complex I in the steady state in this highly longevous
species.

PMID: 9733090, UI: 98401926

----------

Mech Ageing Dev 1998 Jun 15;103(2):133-46


H2O2 production of heart mitochondria and aging rate are slower in
canaries and parakeets than in mice: sites of free radical generation and
mechanisms involved.

Herrero A, Barja G

Department of Animal Biology-II (Animal Physiology), Faculty of Biology,
Complutense University, Madrid, Spain.

Birds have a maximum longevity (MLSP) much higher than mammals of similar
body size in spite of their high metabolic rates. In this study, State 4
and State 3 rates of H2O2 production were lower in canary (MLSP = 24
years) and parakeet (MLSP = 21 years) than in mouse (MLSP = 3.5 years)
heart mitochondria.  Studies using specific inhibitors of the respiratory
chain indicate that free radical generation sites at Complexes I and III
are responsible for these differences.  Main mechanisms lowering H2O2
production in these birds are a low rate of mitochondrial oxygen
consumption in the parakeet and a low mitochondrial free radical leak in
the canary. Strong increases in H2O2 production during active respiration
(State 3) released by addition of ADP to pyruvate/malate-supplemented
mitochondria are avoided in three species because the free radical leak
decreases during the transition from State 4 to State 3 respiration.
These results, together with those previously obtained in pigeons and in
various mammalian species, suggest that the rate of mitochondrial free
radical production correlates better with the rate of aging and the MLSP
than the metabolic rate. They also suggest that a low rate of
mitochondrial H2O2 production is a general characteristic of birds,
animals showing very slow aging rates.

PMID: 9701767, UI: 98367089

----------

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.life-extension,misc.health.alternative
Subject: Re: anti-oxidants and cancer prevention
Date: 18 Feb 1999 00:56:37 GMT

In <7aeois$t7s$1@pegasus.csx.cam.ac.uk> ag24@mole.bio.cam.ac.uk (Aubrey
de Grey) writes:

>The work of Gustavo Barja and coworkers (posted by Steve) is the most
>comprehensive and significant with regard to comparing mitochondrial
>free radical leak: earlier papers, published in 1993 and 1994, are the
>core work.

   But as I seem to recall, they compare organisms of different sizes,
like humans and pigeons.  They're known to have different numbers of
mitochondria per cell (due to the difference in specific metabolic
rate) and so you might expect that each individual mitochondrion would
be tuned differently, since there are more of them for each cell to
deal with.  The radical production rate PER mitochondrion is important
only in comparing critters that have comparable numbers of them per
genome.  Otherwise it tells you nothing.   As I say, I see that this
has now been done, but it's very recent work.


> I'm a bit surprised at Steve's surprise, since he attended
>a talk I gave on this topic only a few months ago!

    I missed your point that this had been done on animals of the
comparable metabolic rates, but different life spans, if you did indeed
make it.  We all agree to await the bat data.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.life-extension,misc.health.alternative
Subject: Re: anti-oxidants and cancer prevention
Date: 19 Feb 1999 11:05:02 GMT

In <36CC1AD4.A36F2E83@servtech.com> Ed Mathes <emathes@servtech.com>
writes:
>
>Steven B. Harris wrote:
>
>> > bat mitos.
>> >
>> >Kate
>>
>>     Holy free radical, Kate, I think you're onto something.  To the bat
>> caves!
>
>Bat mios....is that like guano?  We used that to make crude gunpowder as
>foolish teenagers (saltpeter).


   No, but bat guano is of course the original source of the
unfortunately but aptly named guanine and guanosine.  It's high in DNA
fragments, which bats don't work over much, because like birds they eat
so much flesh they get more dietary DNA than they can use.  As for the
saltpeter, I believe that's made by bacteria in the guano, and in most
other mammals.  Urine test strips these days even test for nitrites,
looking for coliforms.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med
Subject: Re: Where do free radicals come from?
Date: 28 Mar 1999 15:40:26 GMT

In <19990323163529.08823.00000612@ng21.aol.com> pillarinc@aol.com
(Pillarinc) writes:

>A Free Radical is an unstable molecule, often a form of oxygen, that
>react with other molecules in a destructive way. Free radicals are
>produced in the body by fat metabolism,sunlight,radiation,air
>pollution,harmful chemicals,food additives tobacco
>smoke,infection,stress, and numerous other substances.


Comment:
   And produced by the body deliberately and healthfully, to kill
bacteria, signal the need for inflammation and repair, and even to
signal that more blood is needed for some function (when you think hard
about this message, the areas of your brain that use more blood will do
so because of production of nitric oxide, NO, a free radical molecule
which your brain makes and uses to tell blood vessels in a specific
area to dilate).

   Free radical molecules are simply molecules which have one or more
unpaired electrons in a molecular orbital. (A molecular orbital is, for
those who must know, a mathematically defined region of space which the
laws of quantum physics decree that electrons with certain energies and
angular momenta and spin must occupy, on a statistical basis, in a
molecule).  Such molecules are not unstable in and of themselves,
though they are usually quite reactive, since electrons in pairs are
often of lower energy.  But it's all relative.  The oxygen in the air
you breathe is a free radical (it has TWO unpaired electrons).  In
electronically excited oxygen molecules where there are no unpaired
electrons (singlet oxygen), you have an even more reactive molecule,
which would be very bad for you if you breathed it.  So the rule that
free radicals are always the most reactive form of molecules isn't even
true.  We live immersed in free radicals.  We breathe them, and we use
them.  Our bodies make them and sop them up, all quite deliberately.
It's the grossest form of crass pop biology to regard them as simply
"bad things."

                                        Steve Harris, M.D.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med
Subject: Re: Where do free radicals come from?
Date: 29 Mar 1999 09:17:52 GMT

In <36FEF595.F1AE1C15@emory.edu> Andrew Chung <achung@emory.edu>
writes:

>I would also add that there is accumulating research evidence that
>reactive oxygen species (as a class of free radicals) at less than
>cytotoxic concentrations may function as second-messengers (analogous to
>cyclic AMP and GMP).


   Yes indeed.  What better way for a cell to keep tabs on its energy
production than to watch .O2- levels?  The free radical .NO works on
precapillary smooth muscle via a cGMP mediated second messenger system,
as we all know from the Viagra story.  But once upon a time it may have
done some of the job directly. That's a rule in biology:  often you
find complicated systems which show signs of having once worked much
less efficiently, but more directly.  For example, the amino acids
specified by the tRNA genetic code are a little more likely to interact
directly with those DNA codons even without the intermediate protein,
than you'd expect from chance.  So this system may have once opperated
in a very crude fashion even without the "adaptor molecule" which Crick
once so brilliantly predicted would be found (well before it was).  And
some reptiles use pieces of jaw to transmit vibration to their inner
ears, an extremely crude rendering of the beautiful 3-bone system that
developed from parts of the reptilian jaw 200 million years ago in the
first mammals.  Many "second messenger" systems may be the same sort of
thing.  Once upon a time, the free radical did the deed directly.

                                      Steve Harris, M.D.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med
Subject: Re: Where do free radicals come from?
Date: 30 Mar 1999 15:17:25 GMT

In <37075701.12091903@netnews.worldnet.att.net>
Xiuh#tecuh#tli@world#net.att.net (Kolaga Xiuhtecuhtli) writes:


>>"Steven B. Harris" wrote:
>>
>>> In <19990323163529.08823.00000612@ng21.aol.com> pillarinc@aol.com
>>> (Pillarinc) writes:
>>> the areas of your brain that use more blood will do so because of
>>> production of nitric oxide, NO, a free radical molecule which your
>>> brain makes and uses to tell blood vessels in a specific area to
>>> dilate).
>
>I read that Zestril (linospiril?  that ACE-inhibitor) causes the
>production of nitrous oxide which the manufacturer says is of benefit
>to patients with heart failure.

    Nitric oxide (.NO) not nitrous oxide (N2O).  Very different things.
And yes, ACE inhibitors, by increasing bradykinins, do increase NO
production.  How much of their beneficial effect is due to this is
still under study.  Probably comparison with the angiotensin II
blockers (which have many of the same effects but induce far less
bradykinin) will provide some of the answers.


>Anyhow, is there a test to determine the level of free radicals in the
>body?

    No.

> I wonder if free radicals have anything to do with hangovers.

   Maybe.  Nobody really knows.  That pounding in the head is caused by
dilated vessels, and free radicals in the form of NO are probably the
cause of that.  But "ultimate" causes are still in doubt.


>Perhaps free radicals are just like cholesterol and porridge.  It's
>important that the stuff is "just right" as Goldilocks says.


   Exactly.  But "just right" depends on who you are, what you do, etc,
etc.  It's a lifestyle thing as much as anything else.  Are you
planning on being injured or infected?  Or do you want to be free of
arthritis and atherosclerosis in old age?  Do you want enough acid in
your stomach to make ulcers, or are you planning on eating food that
needs the bacteria in it killed?  There are opportunity costs with most
genes.  There are very few that are all bad, or all good.  Just as with
the people they go to make up.

                                    Steve Harris, M.D.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med
Subject: Re: Where do free radicals come from?
Date: 28 Apr 1999 07:59:29 GMT

In <7g4kui$t5u$1@nnrp1.dejanews.com> wdriscol@my-dejanews.com writes:

Harris:

>>Free radical molecules are simply molecules which have one or more
>>unpaired electrons in a molecular orbital. (A molecular orbital is,for
>>those who must know, a mathematically defined region of space which the
>>laws of quantum physics decree that electrons with certain energies and
>>angular momenta and spin must occupy, on a statistical basis, in a
>>molecule). Such molecules are not unstable in and of themselves, though
>>they are usually quite reactive, since electrons in pairs are often of
>>lower energy. But it's all relative. The oxygen in the air you breathe
>>is a free radical (it has TWO unpaired electrons). In electronically
>>excited oxygen molecules where there are no unpaired electrons (singlet
>>oxygen), you have an even more reactive molecule, which would be very
>>bad for you if you breathed it. So the rule that free radicals are
>>always the most reactive form of molecules isn't even true. We live
>>immersed in free radicals. We breathe them, and we use them. Our bodies
>>make them and sop them up, all quite deliberately. It's the grossest
>>form of crass pop biology to regard them as simply "bad things."
>>                                        Steve Harris, M.D.
>
>    It is not correct to say that the oxygen in the air we breath is a
>free radical, since it is found in the form O2, in which all of its
>electrons are paired (some use double bonds).
>                          William Driscoll, Biology student.


    Comment:  William Driscoll, biology student, should become William
Driscoll, chemistry student, for a bit.

     The electrons in molecular oxygen in the ground state (as in the
air you breathe) are not all paired.  Oxygen has the same electronic
structure as nitrogen or carbon monoxide (both triple bonds), but with
two electrons added.  These go singly into two anti-bonding orbitals,
thus cancelling the attraction of one of the three bonds (the three
pairs in bonding orbitals).  The net result is two bonds, but it's
better thought of as three bonds and two half anti-bonds.  If you look
at nitric oxide, which is midway between these two structures
electronically, you'll find a molecule held together with
two-and-a-half bonds (!).  No, I'm not kidding.  It is, of course, a
free radical also.

   It's easy to prove directly that oxygen in the ground state is
triplet not singlet, which is to say it has two unpaired electrons,
instead of all electrons paired up.  Molecules with unpaired orbital
electrons are like atoms with unpaired electrons-- they are
paramagnetic.  If you put them in a magnetic field, they make it
stronger as the unpaired electron spins line up with the field and add
to it (like the iron core in a magnet-- iron atoms have unpaired
electrons also).  Because each atom is a little magnet, the atoms are
atracted in one direction in an inhomgenous magnetic field. So not only
is iron attracted to a magnet, but so is oxygen.  In a chemistry class
you can watch the teacher pour blue boiling liquid oxygen between the
poles of a powerful magnet and watch it hang there, in a boiling liquid
rope running from one pole to the other.  Turn the magnet off, and the
thing collapses to the table, whoosh, like the water-snake alien probe
in _The Abyss_ when they slammed the door on it.  Most impressive.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med
Subject: Re: Where do free radicals come from?
Date: 29 Apr 1999 06:15:24 GMT

In <7g8ise$g9f$1@metro.ucc.usyd.edu.au> jones_j@alf.chem.usyd.edu.au
(Jeff Jones) writes:

>I think that there is a lot of confusion about the term "free radical"
>as it applies to the human body. Obviously there are a lot floating
>around in our bodies, and we are able deal with them given a proper
>diet. It is known that exercise produces a lot more free radicals, but
>presumably your body adapts to deal with this the same as it does with
>a lot of things. It's when you upset the balance too much that you run
>into troubles...
>
>Does anyone know or have a list of compounds that are considered nasty
>free radicals in our bodies? I don't think O2 is one of them.
>
>Jeff


    That depends on your point of view.  Being a biradical, oxygen
accepts free electrons from "radical" oxidants, like Fe(II) atoms (each
of which has unpaired electrons to give up, so selection rules are not
violated).  That is the basis for aerobic life-- the fact that oxygen
is fairly stable toward being an oxidant for other stuff (removing
electrons from things) at body temps, but metal catalized radical
reactions can lower the activation energy enough to do the job at body
temps.  Or room temps.  That makes things nice and controllable, like
an explosive with a well-behaved detonator.  So useful for high energy
output.  If it weren't for O2 and its odd properties, animals could not
exist, and all life would be plants and bacteria and fungi and so on.
You don't get much brain with no battery.

   The fact that O2 triplet is a biradical means that you can't get it
reduced to water without going through mono-radical steps, and those
are nasty.  As you know, your body uses an iron atom in a cytochrome
(a3) to give molecular oxygen an electron in mitochondria.  That
results in a considerably nastier radical, called superoxide: O2.-

    Superoxide attacks double bonds in fats (more about that later) and
also combines with nitric oxide, NO., another free radical signal
molecule, to form one of the nastiest of all oxidants, peroxynitrite
ONOO-.  This stuff chews on DNA.  Superoxide itself is not too bad
since the body has long ago learned to deal with it.  It's used with
Cl- to make HOCl to kill bacteria.  The body also adds more electrons
to it with other enzymes in the mitochondria, to get it oxidized to
water (the protons flowing into this chain drive membrane pumps making
ATP). Any O2.- which escapes is dealt with by dismutating it (where a
pair of superoxide molecules is reacted to reduce one and oxidize the
other) using Mn+2 as a catalyst which gains an electron from one .O2-,
then loses it to the other .O2- (one odd species of bacteria uses
simple Mn+2 still for this).  Nowadays all other life puts Mn2+ in an
enzyme called Mn-SOD (Manganese superoxide dismutase), and converts
O2- + .O2- --> .O2. (normal O2) and O2= (peroxide). Peroxide sucks up
protons to become hydrogen peroxide, H2O2.  This compound is not a
radical, but is nasty as an oxidant, and because of the Fenton reaction
in which H2O2 and free Fe(II) generate .OH, the hydroxyl radical, which
also attacks anything in sight (it's what's produced by radiation, and
is hands down the nastiest free radical of them all in the body).  H2O2
does little if it doesn't find free iron, since your body deals with it
by catalase (fastest enzyme in the body) which converts it to .O2. and
water.  That means H2O2 is not so bad, since organisms long ago learned
to hide free reduced Fe(II) ions, and even organisms without catalase
survive.  (However-- don't underestimate Fe(II): if you bleed into your
CSF from an aneurism in a subarachnoid hemorrhage, you may have nothing
but headache at first, but a day or two later the Fe(II) from heme
breakdown can generate enough radical oxidation to fry your brain).
By contrast with catalase, Mn-SOD is (not surprisingly) required for
aerobic life (unless you're that particular bacteria which uses Mn).
Free radicals from oxygen are that bad.

   In short, .O2. causes problems by what it turns into, and also
because it amplifies other free radical reaction chains when present,
causing oxidative damage in the process.  If a fat molecule with a
system of double bonds in conjugation encounters an odd free radical, a
proton can be abstracted from a secondary carbon due to resonance
stabilization of the resulting carbon centered radical.  Such compounds
react with molecular oxygen to give alkyl hydroperoxide radicals, which
continue to abstract protons from other unsaturated compounds and allow
them to be subject to similar oxidation when they encounter free
oxygen.  In this way, oxygen, due to being a radical available to feed
radical reactions without quenching them, feeds a kind of autocatalysis
where polyunsaturates (particularly w-3 unsaturates like linolenate)
are oxidized via a free radical path.  The resulting molecules form
polymerized gunk and a lot of heat, which is what makes oil based
paints "dry" (oxidize and polymerize via radical pathways) and also the
"drying oils" in them (linseed) spontaneously combust.  The same
happens in your tissues.  There's a trementous burst of free radicals
made when oxygen is re-introduced to tissues damaged by non-perfusion,
and it mostly appears to come from Fenton-type reactions initiating
chains of radical catalized oxidation, driven by molecular O2, and
burning unsaturated fats.  Pretty soon, you're paint thinner.

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med,sci.chem
Subject: Re: Where do free radicals come from?
Date: 30 Apr 1999 05:00:26 GMT

In <3726AFBE.A8EFE44C@cs.uoregon.edu> Bret Wood
<bretwood@cs.uoregon.edu> writes:

>The concept of a "double bond" is a simplification of what's
>really going on.  In actuality, the atomic orbitals of the
>separate atoms mix to form molecular orbitals.  (Orbitals which
>are like atomic orbitals, but are "spread" over the entire
>molecule)
>
>Usually, if you mix two atomic orbitals, you get one unstable
>MO, and one stable MO.  The stable MO is used to form a bond
>between the atoms.  If there are enough electrons to fill BOTH
>MOs, then you end up with no net bond, because the second MO is
>unstable, and cancels the effects of the first one.


Comment:
   Note that this only means unstable or rather "destabilizing"
for the molecule as a whole.  There's nothing "unstable" about
the antibonding orbitals themselves.  They are perfectly
respectable regions of space in which to put electrons, when
regions of space corresponding to electrons of lower energy, and
different angular momentum and spin, are full up.  The only
difference is that bonding orbitals are those where electrons are
more likely to be found *between* nuclei, where their attractive force
holds the nuclei together (that's what a chemical bond is); and the
antibonding orbitals are those outside the space between the nuclei,
where their presence causes the nuclei to be pulled apart:   If "*" is
a an atomic nucleus and "." is a place where bonding electrons are
often found:

                          .......
 antibonding region)*  bonding region *(antibonding region
                          .......


>in oxygen, it turns out that the way the two oxygen atoms'
>atomic orbitals combine, there is a pair of molecular orbitals
>of equal energy.  And it also turns out that there are two
>electrons to fill these two orbitals.


Well, it takes four electrons to fill these 2 equal energy pi
antibonding orbitals (one pair for each), but the two in ground
state oxygen go one per orbital (triplet state).  Each of them
spends most of its time in a region away from the region between
the nuclei.  The first electron (O2+) actually goes into BOTH of
antibonds, since they are equivalent.  It has to go into an
antibond, since the space between the oxygen atoms is "filled up"
after the molecule has 14 electrons (the N2 or CO or structure).
By filled up, we mean that it's possible to put electrons in
there, but they are so far from the nuclei that antibonding
molecular orbitals fill up first, and pull the molecule apart.

These 14 electrons in N2 or CO go to fill seven orbitals (2
electrons in each orbital, spinning in opposite directions):  2
pi bonds, 3 sigma bonds, 2 sigma antibonds. The total bond in
these molecules (N2 and CO), thought of as triple (5 bonds minus
2 antibonds) are among the strongest known molecular bonds (I
believe the bond in CO actually is the strongest known).  In N2
and CO there are roughly 10 electrons in the space between the
atoms (this fills this space up in a molecule this size) and 4
more in sigma antibonds in the space away from both, projecting
out in a hemisphere-like nubbin along the axis running between atoms in
the molecules, but not between the atoms, rather pointing toward the
space outside:

  ()N-=N()<-- sigma antibond electrons

Here -= is the triple bond (hard to draw in ASCII).  I won't
draw the sigma* electrons in what follows (no room) but the reader
should understand that 4 of them are there in all the molecules
discussed. (In practice the 4 innermost electrons, don't
participate in this scheme much, but since they are arranged in
1s2 atomic orbitals spherically about their "parent" nuclei, the
fraction of electrons between the atoms vs. not-between is not
changed by this non participation in molecular orbitals).

That triple bond in N2, BTW, is why it's so hard for living
organisms to fix nitrogen, which involves breaking that bond,.
It's also why there wasn't an industrial process for it until
 WW I, and why the Germans needed the nitrogen for many high
explosives, in which energy is made by re-forming N2.  Fixed
nitrogen was the "atomic bomb" of WW I (a way to densely store
coal and oil energy compactly for sudden release against the
enemy), and it and some iron mines in France kept Germany in the
war long after it should have been out.

I can write triplet molecular oxygen, which has 16 electrons, as

 .O=O.

showing only the 2 equivalent pi antibond (2p pi*)
electrons outside of the N2 structure.  These are each equally
present at each atom, and though there is one per atom, they
should not be thought of as going in "one on each atom."  They
are each smeared out in a pi antibond orbital which includes four
lobes which point away from the oxygens, out into space (total of
4 lobes for each oxygen).  However, the above structure does
suggest some of the truth, which is that these electrons spend
most of their time away from the space between the atoms.

With regard to the N2 structure above, the "extra" 2 electrons in
molecular oxygen don't both like to go jammed into the same flat
orbital.  This is basically because they can be farther apart
(use more space) if they use the two different but energetically
equivalent ones available (the energy to jam electrons into the
same orbital is about 40 kcal/mole, which is roughly half of the
total possible energy given up in forming a chemical bond in this
type of molecule and a pair of bonding electrons).  Such
electrons in two "pi antibond" (pi* or pi-star) orbitals have
their total angular momenta (a vector sum of molecular orbital
angular momenta and intrinsic electron spin) pointing partly in
the same direction, and since the vector sum of this is what
determines magnetic properties in an atom or molecule, oxygen
molecules have a net magnetic field in the triplet or ground
state.

If one electron is moved and flipped in spin to the other orbital
so the two electrons are paired and the molecule is flattened by
having electrons in only one antibond orbital (singlet oxygen),
the magnetic field goes away, as the all electrons are paired and
effects of all motions cancel. You can think of the antibond
electrons in this molecule as like two elliptical loops of wire
on top of each other with current running opposite directions.
It has flat (planar) symmetry, and due to the opposite currents
there's no net magnetic field.  Tip one loop up at 90 degrees,
however, and the resulting sausage has some axial symmetry, and
now has a net field, since the two magnetic fields from each loop
add at an angle of 90 degrees.  That's what happens when singlet
oxygen falls to ground state triplet oxygen.

If you add one more electron to molecular oxygen you get super-
oxide molecular ion, O2-, which looks like

   - :O=O.             or            .O=O: -

Here the "-" indicates a net negative charge on the molecule.
(You might guess that the negative charge isn't on one end or the
other-- it's spread over the molecule with a little more at each
of the ends, which are equal.  The reality is halfway between, or
a mix, of the structures above).  And the bond is further
weakened by half a bond, so these atoms are held together now by
only 1.5 "bonds" (even though still symbolized as "=" ).  But the
molecular ion is still stable in vacuum and still magnetic.  In
the body it reacts with water, but not nearly as quickly as many
other free radicals do (in this it is something like nitric
oxide).  There are solid compounds containing superoxide ions
such as above in a crystal lattice, and superoxide salts of
metals are what forms when you burn the larger radius alkali
metals in lots of oxygen: for example: KO2, which is a salt of K+
and O2.- ions.

If you add one more electron to superoxide you get peroxide,
which I can write as

  - :O--O: -

Note that we're now down to one bond (and the atoms move apart
from 1.28 to 1.49 A in consequence) and of course we now have two
net negative charges on the molecular ion.  Metal peroxides are
known as well (sodium when burned forms lots of Na2O2), and you
know that the corresponding Lewis acid is hydrogen peroxide.
This is nasty stuff in the body also, but not nearly as nasty as
superoxide.

If you subtract an electron from an oxygen molecule you get

  .O=O+    or     +O=O.

where the positive charge is equally distributed (1/2 a charge on
each atom) and the bond is 2.5 bonds-- now stronger than
oxygen's.  O2+ ions are stable in vacuum, but obviously pick up
electrons from whatever they can (anything except each other). It
reacts almost instantly with water.  It has the same electronic
structure as nitric oxide .N=O with just one lone electron in a
2p pi* state, again equally present at both atoms in O2.  The
single electron (again) is actually distributed in two pi
antibonds, each of which has 4 lobes sticking out into space at
angles in a way that is hard to draw, but looks like  ) O=O (
where the curved things are four lobes bent away at angles like 4
petals of an opening four-leaf clover flower, or a dish (which I
can only show here in cross section).

")" and "("  signify here something less than half an electron
(average), which is to say, a bit less than 1/2 the density of
the one pi antibond electron:


 ) O=O (                 O2+  (1/2 positive charge at each end)

 ) N=O (                  NO

The NO (also can be thought of as .NO) is slightly polar, even
though uncharged as a whole.  It's a signal molecule in the body,
and chemically is a stable gas that can be bought in a tank, and
is used medically in very low concentrations (diluted in
nitrogen) as a bronchodilator (it is highly poisonous in larger
doses than few parts per million, and is not to be confused with
the anaesthetic nitrous oxide N2O, or laughing gas).   O2+, like
O2- is charged but nonpolar.  O2+, O2, O2-, and NO are free
radicals.

The two lowest electronic energy states of molecular oxygen are
called triplet and singlet (spectroscopic terms denoting states
for two electron systems which have unpaired or paired electron
spins).  Ordinary oxygen in the air is in the triplet state.  One
can write the higher energy singlet O2 (where both antibond
electrons go into the same orbital) like this:

  > O=O <    Singlet oxygen (semi-planar)

where "<" this time represents about one electron (again each of
the two electrons is at both ends of the molecule and isn't
localized). Now the two electrons are more nearly stuck in the
same plane (think of them as in four lobes total, two coming from
the direction of each atom, and all lobes lying along axes within
a plane flat on the screen).  You can see that the electrons are
more closely confined, and it's not hard to see why they don't
"like" it.

Normal ground state triplet oxygen is our thus our O2+ and O2-
structure again, and would be drawn with 1/2 of the lobes of the
pi antibond orbitals at each end twisted at 90 degrees along the
axis connecting the oxygen atoms, so two lobes are in the screen
plane "<" and two are not in the screen plane "-".  This is what
I've been drawing in cross section as:

  )-O=O-(      triplet oxygen (non-planar)

"-(" is the opening-four-leaf-clover-like affair
previously described.  This is difficult to draw in this medium,
and I hope readers can visualize it.  A better description would
be four pear-shaped balloons all stuck small end downward into a
pot just big enough to hold the ends together.

>One interesting property of oxygen which is a direct result of
>these two unpaired electrons is that "triplet" oxygen (oxygen
>where the unpaired electrons have the same quantum spin) has
>a different chemical reactivity than "singlet" oxygen in many
>types of reactions.

Yes.  A molecule with two unpaired electrons naturally behaves
differently from one with a paired set.

>I'm not sure that I've ever actually heard the term "free
>radical" used to describe molecular O2 before though.  But it
>does fit the definition, even though it isn't as reactive as
>most free radicals.

No.  However, free radicals differ greatly in reactivity.  There
are, for instance, a great number of nitroxide "spin trap"
molecules which are stable free radicals (like NO.), but fairly
unreactive solids (many can be given as pharmaceuticals, like
PBN).  They are paramagnetic and all feature the odd group

  \
  .N=O
  /


>Most free radicals are reactive because electrons desperately
>"want" to pair up in molecular orbitals.  I'm not exactly
>certain why this isn't the case in O2, but I believe it is
>because addition of an extra electron causes the orbitals to
>become asymmetric, and _then_ that extra electron ends up
>being in an "antibonding" orbital, which is not a good thing.

The two highest energy electrons in the two lowest energy states
of O2 are in antibonding orbitals, whether paired or not.

Electrons want to pair up only so that they can share the same
physical space (orbital) when space in low energy conditions is
at a premium (as often happens between atoms in chemical bonds).
However, pairing brings an energy cost, since the electrons are
nearer each other when in the same orbital than when not, and
they still electromagnetically repel more than they magnetically
attract.  If that energy debt isn't paid for by being able to put
another electron in the lower energy orbital, it won't happen.
If you look at transition metals, you see many examples of
electrons spreading out in equivalent orbitals and trying not to
pair.

Electrons occur in pairs so often in chemistry because light
elements are far more common and so are elements with even
numbers of protons (due to helium fusion and proton pairing
energy forcing nucleosynthesis in certain directions in
supernovas.)  It takes an odd-proton element in a molecule other
than hydrogen to give the molecule an odd number of electrons,
and many of these molecules when formed from light elements
simply react with each other.  Even if this happens slowly it
uses them all up and this depletes light element molecular
radicals in nature. If it were not for sun energy (including life
which uses it) there wouldn't by any light-element molecular
radicals, and this includes molecular oxygen, as well (a very
unnatural thing at 300 Kelvin, and surely not to be found in
quantity on lifeless worlds at those temps).  Among larger atoms
(transition metals) compounds containing free radicals
(paramagnetic compounds) are the rule on Earth.  But those elements are
far less common.  It's not a coincidence that of those which are
most common, our bodies use nearly all as catalysts for various
reactions, including those which use and transport O2.

>Similarly, if you remove an electron, then the orbitals become
>symmetric, and there is a "bonding" orbital with only one
>electron in it.  That is also a bad thing.  So, although the
>electronic structure of O2 is weird, it's a special situation
>where nature seems to make the best of an unusual circumstance.

The above assumes the wrong electronic structure for oxygen.
Anyway, read the tutorial.   If one removes an electron from
oxygen you get a molecule with only one antibonding 2pi electron,
very much like NO.  It's asymmetric as can be.

NO, BTW, really is an odd molecule.  For example, the 2pi* and
2sigma* orbitals are close enough that NO is the only gas at room
temperatures for which electronic transitions play a significant
role in the heat capacity.  For most molecules (exceptions are
large halogens) even the non-fundamental vibration modes have
frozen out by the time you get to room temp.  The other odd thing
about NO (if you weren't looking at molecular orbitals) is that
singly ionized NO+ has a bond stronger than NO's.  It's a triple
bond stronger than N2's, and 97% as strong as that in CO, and
second in strength only to it.

I'm crossposting it to sci.chem in case I've said something which
needs correction, and also to invite comments.

Steve Harris

From: rparson@spot.Colorado.edu
Newsgroups: sci.med.nutrition,sci.med,sci.chem
Subject: Bonding (was Re: Where do free radicals come from?)
Date: Mon, 10 May 1999 00:18:47 GMT

In article <7gb7lc$a5l@sjx-ixn1.ix.netcom.com>,
  sbharris@ix.netcom.com(Steven B. Harris) wrote:
> In <3726AFBE.A8EFE44C@cs.uoregon.edu> Bret Wood
> <bretwood@cs.uoregon.edu> writes:
>
> >The concept of a "double bond" is a simplification of what's
> >really going on.  In actuality, the atomic orbitals of the
> >separate atoms mix to form molecular orbitals.  (Orbitals which
> >are like atomic orbitals, but are "spread" over the entire
> >molecule)
>
> >Usually, if you mix two atomic orbitals, you get one unstable
> >MO, and one stable MO.  The stable MO is used to form a bond
> >between the atoms.  If there are enough electrons to fill BOTH
> >MOs, then you end up with no net bond, because the second MO is
> >unstable, and cancels the effects of the first one.

 [A great deal of reasonable stuff deleted.]

> I'm crossposting it to sci.chem in case I've said something which
> needs correction, and also to invite comments.

 A comment not on Steve's post per se, but on the molecular orbital
 concept in general (triggered by the first sentence of Bret Wood's
 post .) The whole Molecular orbital approach is an approximation,
 and a rather crude one at that. There is a widespread misconception
 that molecular orbital theory is superior to valence bond theory. It
 it not - it is both quantitatively and qualitatively inferior.
 The simplest MO treatment gives an H_2 binding energy of 0.099
 hartrees, the simplest VB gives 0.116 hartrees, and experiment gives
 0.167 hartree. Worse, the MO description turns into complete nonsense
 when you try to describe molecular dissociation - all the way up to
 the SCF level, MO thinks that H2 dissociates into a completely
 nonphysical 50-50 superposition of neutral atoms and H+/H- ion pairs.
 VB correctly dissociates H2 into neutral atoms. MO does, of course, get
 some things right that VB doesn't, such as the paramagnetism of O2, and
 it is a more useful starting point for understanding molecular
 electronic spectra. But for a basic qualitative accounting of most
 chemical facts VB is superior.

 As you fix up the simple theories they must converge towards each
 other, and it is a standard P. Chem. exercise to show that adding
 ionic terms to the VB wavefunction for H2 is mathematically equivalent
 to adding configuration interaction to the MO wavefunction. In general,
 MO provides a more computationally _convenient_ starting point for
 a sophisticated calculation, so most such calculations are  phrased
 in MO language. (Not all, though - W. Goddard and his students have
 gone a long way with Generalized Valence Bond theory.)

 Historically, VB was taken up first by non-physical chemists, so
 when MO-based ideas were later used to explain phenomena that VB
 couldn't, some people got the idea that MO was a more inclusive
 theory of chemical bonding. It isn't, it's a complementary theory.

 ------
 Robert

From: sbharris@ix.netcom.com(Steven B. Harris)
Newsgroups: sci.med.nutrition,sci.med,sci.chem
Subject: Re: Bonding (was Re: Where do free radicals come from?)
Date: 10 May 1999 07:22:54 GMT

In <37366526.9BDB05F0@cs.uoregon.edu> Bret Wood
<bretwood@cs.uoregon.edu> writes:

>Indeed.  VB has a more developed qualitative approach, and MO has
>a more developed quantitative approach, but both models can be used
>to describe any system.  But I wouldn't know where to being describing
>the paramagnetic nature of O2 using a qualitative VB description.
>
>Is it just that the single-bonded diradical resonance structure is
>much lower in energy than is normally the case for a double bond?
>Is there any particular reason that this resonance strucutre is
>more stable than is typical?  You usually don't see a lot of
>contribution by diradical analogues of double bonds in most organic
>structures, as far as I know.  (Except for spin traps obviously)


   The interesting resonances are not
                     .     .
  :O  :: O:   <-->  :O  :  O:
   ..    ..          ..    ..


   For the first predicts paired electrons and the last a single bond
(resulting in 1.5 bonds strength, when experimentally it is two.

   What you want is:

    .                               .
 - :O  ::  O: +    <-->   + :O  ::  O: -
    ..     .                 .      ..



           . ___ .                   . ___  .
    or    -O <-- O+       <->       +O -->  O-



These are obviously much better than VB structures which have only one
bond, and mix of them causes no net charge separation or polarity.

From: Steve Harris <sbharris@ix.netcom.com>
Newsgroups: sci.med.nutrition
Subject: Re: Chromium - wipes me out
Date: 14 Jun 2005 09:30:24 -0700
Message-ID: <1118766624.303704.254150@z14g2000cwz.googlegroups.com>

>>Just remeber one thing:  free radical DAMAGE is the key to "chronic
disease," and if you search on pubmed.com you'll find many researcher
saying this explicitly.<<


COMMENT:

They can say it as explicitly as they want, and that doesn't make it
generally true. It's a theory that goes back to Harman and has been
promoted by optomists like Pearson and Shaw, and has gained currency as
a good hypothesis. The problem is that there's not much evidence for
it.

Now wait, you say. Whenever there's worn out or damaged tissue, aren't
there more free radicals around?  Sure. Whenever there's damage to
buildings by fire or crime, there are lots of cops and firemen around.
This does not mean that uniformed people or paperwork or insurance
claims cause disasters.  Try *removing* the cops and firemen to test
this hypothesis. It doesn't work.

Nor does removing free radicals modify chronic disease progression,
much. The biggest cause of chronic disease in developed countries is
aging. If you remove free radicals by introducing spin traps,
antioxidants, and so on, you do not find aging markedly changed in
animals. There were some hopeful experiments along that line, but they
haven't been replicated, and some of the early ones turned out to be
due to dietary restriction, not antioxidants (these things often taste
bad).  Nor is there much evidence that you can modify the rate of
"chronic disease" in humans by giving antioxidants. I've brought up the
subjec of vitamin E, and all the proponents can think of to say is that
maybe we have to use every kind of molecule with vitamin E activity
that exists in nature. Well, they didn't say that 25 years ago before
the vitamin E experiments failed. The claims of the vitamin hawkers are
moving targets. Produce data against them, and they just change.  They
are not only data-free, but data-proof. When it's explained to these
people that the greatest life expectancies and lowest chronic disease
rates occur in the Japanese, who probably eat one of the highest
polyunsaturated fat diets there is, they just blank out and invent
magical reasons that the Japanese might be eating huge doses of free
radical blockers to make up for it.

You wouldn't want to just damp out all free radicals willy nilly
anyway, even if you could. They are the body's messengers for dealing
with infection, and for triggering healing. Yes, sometimes they need
modulation. Free radicals are produced in copious amounts during tissue
damage (acute trauma or infection or tissue ischemia) and they often
are produced in amounts too large for optimal health (due to the fact
that evolution did not design you with antibiotics in mind, nor to
survive giant traumas that require ventilators and pressor support or a
lot of brain or heart ischemia--- you're expected to DIE in those
cases). So antioxidants (ironically) actually shine more as possible
modifiers of ACUTE disease processes, not chronic ones. For chronic
disease, the inflammatory process will probably need to be modified far
higher up in the process than the end-result where free-radicals are
generated. Another story.

SBH