Hi,
Im doing a science fair question, and im troubled with my results. My
experiment is does bacteria effect the rusting rate of iron? What I
did was create a varible and a non varible. In the non varible i put a
jar of water with a nedle and bleach. In the other it was the same
thing just with out the bleach. I observed both of the jars and the
one with bleached took an estimate of 60 min longer to rust then the
one with out bleach. I repeated this experiment sevreal times and came
up with the same results.
I asked my teacher and she hadent have a clue so she recommended this
site.
so i want to know why the bleach slows the process of rust? Can it be
because of the killing of bacteria? or just another reason.
PLEASEEE ANSWER BEFORE FEBRUARY 16, 2003. IM BEGGING 
                   ------ THANK YOU IN ADVANCED ------

---

Request for Question Clarification by sublime1-ga on 02 Feb 2003 19:57 PST

webdevil89...

What I found indicated that the bleach should speed the process:
http://newton.dep.anl.gov:70/askasci/chem00/chem00345.htm

Rust is the result of the oxidation of iron, and bleach is a
good oxidizer, which should, theoretically, speed the process,
and the questioner on the page above seems to have found that
it does so.

What metal is the needle made of? You might try using iron
filings from a machine shop.


As for bacteria:

"Scientists have been debating for a century whether iron oxide,
 commonly known as rust, results from biological or inorganic
 processes. The question of whether life forms such as bacteria
 or viruses play a role in the formation of rocks rich in iron
 oxide has not yet been settled, Folk said."

and

"Two geology professors at The University of Texas at Austin,
 Dr. Robert L. Folk and Dr. Kitty L. Milliken, have demonstrated
 that iron oxide filaments from a variety of geological periods
 on Earth are lifelike in form at microscopic levels. They say
 their research could have implications for Martian exploration
 and the search for some form of life on other planets."

"...Folk and Milliken say the microscopic shapes they found would
 strongly suggest that living matter is intimately involved in the
 process."

"Folk is a Dave P.Carlton Centennial Professor Emeritus in Geology
 in the department, who has done pioneering work on forms called
 nannobacteria. Nannobacteria are incredibly small strands, along
 with spherical and egg-shaped objects found in rocks and minerals
 and believed by a few scientists to be dwarf forms of bacteria.
 They are about 1,000 times smaller than normal bacteria. Folk
 said half a billion nannobacterial cells could fit on a pinhead."

"Many biologists say no living creature could be smaller than
 about 0.2 microns because they say that is too small a size to
 contain the genetic material necessary for life. (A micron is
 one-millionth of a meter in the metric system.) And other
 scientists deny the structures represent the presence of life,
 arguing that the suggestive shapes are merely the result of
 chemical actions or weathering."

"Folk and other researchers admit that suspiciously lifelike
 shape does not necessarily prove or disprove the presence of
 life. But they believe the forms are fossils of the most
 primitive and earliest life forms on found on Earth and beyond."
http://www.utexas.edu/admin/opa/news/00newsreleases/nr_200003/nr_bacteria000320.html

So maybe you're on to something there!

A gallery of nannobacteria photographs are available here:
http://www.msstate.edu/dept/geosciences/4site/nannobacteria.htm

sublime1-ga

In order for Iron to rust (in other words to convert from Fe metal to
an Fe+++ oxide or hydroxide) it needs to shed electrons - in other
words to find an electron acceptor. The half reaction is:

Fe --->  Fe+++   +  3e-

The breakdown of hypochlorite (OCl-, or bleach) assists the above
reaction, as follows:

2e-  +  OCl-  +  H2O  --->  2OH-  +  Cl-

If you balance the two equations (which are called "redox", or
reduction-oxidation reactions) you obtain the following:

3OCl-  + 2Fe  +  3H2O  --->  2Fe(OH)3  +  3 Cl-

Notice that the charges on the left hand side of the equation are the
same as on the right hand side... in other words the equation
balances... but the iron has donated electrons to hypochlorite in
order to break the bond between oxygen and chlorine. (Your teacher can
explain this better, and you probably will not get into it in detail
until first year college chemistry).

If you measured the mass of the needle with a very accurate scale I
bet you would find that it had lost mass... i.e. that it had partially
dissolved. Yet you saw no rust - why?

Your first clue can be found at
http://www.adbio.com/science/analysis/ph_scale.htm
You will find that the pH of the bleach solution is quite high - 12.6

Your second clue can be found at
http://www.waterspecialists.biz/html/precipitation_by_ph_.html
Quote from this site:
"Metal hydroxides are amphoteric, i.e., they are increasingly soluble
at both low and high pH, and the point of minimum solubility (optimum
pH for precipitation) occurs at a different pH value for every metal."

Your third clue: 
http://yarchive.net/metal/rust_remove.html
Chloride forms very stable (soluble) complexes with trivalent iron
(and is therefore a component of rust removal compounds)

My theory... you are dissolving your iron more rapidly in the
hypochlorite solution, but are simultaneously forming soluble
coordinate compounds such as iron chloride complexes or iron hydroxide
complexes - so you don't "see" the rust. Your results probably had
nothing to do with bacteria.

The more technical version... Iron(III) tends to form 6-coordinate
compounds, which can include in the six sites chloride, hydroxide or
water molecules. If more than three of these six sites contain
chloride or hydroxide, the species will be a negatively charged ion in
solution. Even in the absence of chloride high pH solutions tend to
form hydroxide-dominant coordinate compounds, such as Fe(OH)6 with a
charge of -3.

For your interest, I used to work in water treatment where we would
remove heavy metals by precipitation at high pH. We had to be very
careful not to overshoot the pH, because if we did the metals
(including iron) would not precipitate - they would actually
re-solubilize. This area of chemistry (coordinate chemistry) is one of
the lesser appreciated areas, but is extremely important in
environmental industries.

Hope this helps, and congrats on your participation in the science
fair. I had the pleasure of being a judge in one such fair a number of
years ago, and came out of it with a great degree of enthusiasm. There
were some exhibits where moms and dads had obviously done 90% of the
work, some exhibits by young show-people who tried to spend their way
to the top with professional apparatus and posters, and then there was
my favorite - the handful of born engineers and scientists who had
thoughtful enough minds to be "troubled by their results" and search
out the answers. Best of luck to you.

Found a good site for you - (I'm a little more awake now.. last
comment was at 3:00 in the morning).

http://www.civil.nwu.edu/ehe/courses/ce-367/Chapters/Lecture4.html

This is a lecture on the solubility of iron hydroxide as a function of
pH, including phase diagrams. Basically it confirms what I suspected,
although the coordinate complex described for iron is a 4-coordinate
(Fe(OH)4)-. (Iron can form 4, 6, or 8-coordinate complexes)

By the way... found your project idea described in a six year old post
at http://www.eskimo.com/~billb/scifair/chem.html, and can add a
couple of comments to the concept. The post:

"BIO-RUSTING"
I heard someone theorizing that the normal rusting process of iron
might depend on bacterial action. If bacteria are not present, will
iron rust extremely slowly? It shouldn't be too hard to see if this is
true. Get two glass jars and two iron nails. Clean the iron with
sandpaper to get rid of any oily coating. Fill the jars with distilled
water. Drop in the iron. One jar should have normal bacteria from the
environment, so add a small bit of dirt. Put a tiny bit of powerful
bactericide in the other jar. (Or sterilize it in the same way you do
"canning" at home) Wait a few days, and see if one piece of iron
obviously rusts more than the other. What kind of bactericide to use?
I dunno, maybe anti-bacterial hand soap. Better put some normal soap
in the other jar as a control, since I don't know if soap affects
rusting too. Nobody has tried this project yet, as far as I know. If
it doesn't work, it still is interesting: it tells you that bacteria
doesn't accelerate rusting. Search for: IRON BACTERIA, IRON-LOVING
BACTERIA, NANNOBACTERIA. (I recently heard that those big chunks of
rust that cover the sunken ship Titanic are full of iron bacteria.)
(Name deleted by Blanketpower)
Original post - Wednesday, September 18, 1996 at 23:13:49 (PDT)

Were you using distilled water? If you were using any form of drinking
water from a municipal source (or bottled mineral water) it was
probably already sterilized by chlorine addition. Did you remember to
add the soil? You have to "seed" the reaction by adding something that
already has bacteria in it.

If want more info on the types of bacteria that are active in iron
oxidation take a look at:

http://www.esr.pdx.edu/pub/biology/limnology/limn-12.htm

The general rule is that virtually any reaction that is exothermic
under any given set of conditions can be "catalyzed" by bacteria.
Bacteria do not try to work against nature... they just look at
reactions that are going to happen anyway, and go "Mmmm... tasty".
They speed up the reaction and on the way they pirate a bit of the
released energy for their own metabolisms. Thus, it would not be
surprising to find such bacteria in rusted hulls. They do not "cause"
the rust... they just sort of come along for the ride in environments
where iron oxidation is already favorable.

In mining we commercialize iron and sulfur oxidation reactions with
one bacteria called thiobacillus ferrooxidans.
(http://www.mines.edu/fs_home/jhoran/ch126/thiobaci.htm). If you are
looking for a way to perform a "fast" demonstration of biological
activity one of the best ways is not to use metallic iron, but rather
to get a piece of pyrite (FeS2) (or you could even use laboratory
grade FeS powder, which the school probably has, although natural FeS2
is better). It's even better if you can mix in some copper sulfide
chalcocite, or Cu2S, because the water will start to turn blue as the
oxidation proceeds, for nice visual effect.

Pulverize the solids in a mortar and pestle, divide the material in
two, and put the two samples in glass tubes with glass wool at the
bottom (to keep the sample from pouring out of the bottom of the
tube). You want a column of iron sulfide sand maybe 4" high and 3/4"
wide, held in by glass wool and open to the atmosphere at the top. Now
get three liters of rain water, river water, puddle water, or any
other kind of water that is not sterile. (Try to avoid getting water
full of road salt if you live in one of those nasty cold areas). If
you could get some water out of an old rusty pipe or a junk yard it
would be ideal, because it would aready have a population including
the right kind of bacteria.

Carefully adjust the pH of the water to 4 using hydrochloric acid
(otherwise you'll spend about a week waiting for the reaction to
start). Now divide your water into three one-liter samples. One will
be used for the bacteria experiment, one for the sterile experiment,
and one liter for make-up water (since you will have some
evaporation). Add a couple of drops of 1.0M AgNO3 (silver nitrate) to
the water for your sterile experiment. (The stuff is like the kiss of
death for T.ferrooxidans bacteria). Pour your "natural" water through
one tube of iron sulfide and your "sterilized" water through the
other. For the next week you will pour the water through the tubes
twice a day, enough to keep the iron sulfide moist. (The reaction goes
on at the moist surface).

Every time you go to pour the water through the tube, record the pH of
the water draining out. The pH is a much more sensitive indicator of
the rate of oxidation, and will go down to about 1.8 as the reaction
speeds up. If everything is going well you should see both changes to
the pH and to the sample appearance of the "non-sterile" sample within
about four days, and quite a big change within 8 days.

I ran a bigger scale of a reactor similar to this a few years ago when
I was cultivating some bacteria for bioleaching work. I had a bit more
of a sophisticated apparatus (for example a pump that let the sample
irrigate for 30 minutes then "rest" for 30 minutes... you would only
be irrigating twice a day). The simple apparatus, though, should work.
Try to keep it in a warm place (it likes abount 90 degrees, and slows
down below 75 or above 110) and avoid putting it under bright
UV-generating fluorescent lights. (Bacteria are shy and like their
privacy).

We estimate in the mining industry that about 95% of the oxidation
responsible for acid mine drainage etc. is mediated by bacteria... in
other words that the oxidation reactions happen at about 20X the rate
that we would see if mines were "sterile". In mines it's a big
problem, since you need to dump the water that you pump out (hopefully
after hydroxide precipitation of the metals)... but sometimes in the
copper industry we put the reactions to work in a process called "heap
leaching" or "dump leaching". Basically you pile the rock on a big
waterproof pad, irrigate it in a manner similar to what I described,
and collect your metals from the blue water that runs off. When you do
this it's all environmentally contained, and you harness what would
otherwise be an environmentally annoying reaction for productive
purposes.

This questions illustrates an important point regarding science
experiments.  You need to be sure that the experiment devised really
tests the hypothesis you want to test and that the control is truly an
appropriate control.  You don't mention your water source or whether
any additional material was added as a source of bacteria, so I am
assuming you were relying on natural bacteria in the environment as
the source.  If you used bottled or tap water (chlorinated), and clean
containers, there was probably fairly little in the way of bacteria in
either container, even before adding the bleach.  The major variable
added was the bleach, a potent chemical which affects many chemical
reactions.  Thus, what this experiment really measured was the effect
of bleach on rusting, not bacteria, which unless added, were probably
in too small number to have as great an effect as the bleach itself. 
(Note this is another hypothesis, which would require a separate
experiment to prove.)
  In general a control (or non-variable) for an experiment should be
as close as possible to the experimental, with the exception of the
variable you are testing.  In your case, the difference between the
two was addition of bleach to one jar, otherwise the two were
identical. This makes the addition of bleach the variable.  Then
differences between the two need to be explained in terms of effects
the bleach  may have caused, which could relate to either killing the
small amt of bacteria present or as a direct effect of the bleach on
the chemical reaction.
   A better way to have set this up would be to start with identical,
sterile set-ups (sterile containers filled with sterile, distilled
water, each with a sterile needle, all sterilized through non-chemical
means, so as not to leave behind anything which would kill the
bacteria).  Then add bacteria to the experimental container, but not
the control.  It would be best to add as pure a form of bacteria as
possible to avoid other confounding variables (such as chemicals in
the soil, water, etc that the bacteria came from).  Understandably,
this is difficult, so the best one can usually do for a school project
is to add something known to contain a lot of bacteria.  Of course one
could add this same substance to the control after first sterilizing
it through non-chemical means as well, which would improve the
control.