Salt and vinegar for rust removal

On Sat, 8 May 2004 06:02:31 -0700, "Paul O." <[email protected]>
posted:

>Did a google for rust removal and saw a few references for removal of
>lightly rusted hand tools using table salt and vinegar. Any of you use this
>method? If this works would like to try it before scrounging up the parts
>for electrolytic rust removal. How much vinegar and salt do you mix with
>water? Thanks.

I'm wondering why the salt. It surely does nothing but perhaps cause
problems later on. I would prefer a stiff brush and kerosene or WD40.
Although vinegar will not rapidly dissolve iron, it will very slowly.
It will also, a little more quickly, dissolve rust, but I would prefer
to remove this mechanically with something that impedes further
rusting. Just my 3 cents...

In article <[email protected]>,
[email protected] says...
> >Did a google for rust removal and saw a few references for removal of
> >lightly rusted hand tools using table salt and vinegar. Any of you use this
> >method? If this works would like to try it before scrounging up the parts
> >for electrolytic rust removal. How much vinegar and salt do you mix with
> >water? Thanks.
>
> I've used sandpaper, then fine steel wool. Works just fine, and gets
> rid of heavy to light rust.
>
And completely destroys any collectible value the tool may have
had.

--
Where ARE those Iraqi WMDs?

In article
<[email protected]>,
[email protected] says...
> I have used the salt and vinegar method for cleaning small hand tools for
> several years. It works slowly, but works well for high carbon steel --
> especially plane irons and chisels. You will need to use something slightly
> abrasive to help clean the surface --

I've also used it many times, but I never needed an abrasive.
Wipe it off afterwards or maybe use a soft toothbrush, but
that's all.

> The mixture is simple, standard 5% acidity white vinegar and table salt --
> just dissolve as much salt in the vinegar as it will take, and it takes
> quite a bit and dissolves slowly as it nears the saturation point. The
> mixture can be reused several times even though it turns red from the rust.

My experience is the same.

> Another important point, the steel will rust very quickly when removed from
> the solution. Keep some clear rinse water and a can of WD40 spray handy to
> clean and protect the tool.
>
Yea, verily, amen! Rust starts reappearing within a few minutes
:-).

--
Where ARE those Iraqi WMDs?

In article <[email protected]>, [email protected]
says...
>
> I know nothing of chemistry so my questions are(1) When disposing of this
> stuff, anything in particular to worry about? (2) When storing, is a metal
> coffee can ok, should the container be vented? Don't want to use this
> mixture in anything that it will eventually eat thru. Thanks all for your
> help and advice.
>
I'm no chemist either, but a friend who is told me the salt
(sodium chloride) and the vinegar (acetic acid) reacted to make
a weak hydrochloric acid. So I wouldn't store it in metal.

I have some in the original vinegar bottle (plastic) that's been
there for at least a couple of years. As someone pointed out,
it gets dirty but it still works. I do add a little more salt
from time to time.

--
Where ARE those Iraqi WMDs?

In article <[email protected]>, Sandy
<> says...
> I'm wondering why the salt. It surely does nothing but perhaps cause
> problems later on.
>
See your friendly local chemistry professor :-).

--
Where ARE those Iraqi WMDs?

In article <[email protected]>,
[email protected] says...
> Ditto! All I can say is that the presence of ions in solution, ionic
> strength, does definitely affect how species in solution react. Maybe there
> is some physical chemistry website or ng you can visit and ask this
> question. I'd be interested to know, too!
>
Wow! I didn't mean to start such a learned discussion :-). My
knowledge of chemistry is limited to making various explosive
compounds, learned long ago in my juvenile days. And lately, I
think I've forgotten most of that - CRS seting in :-).

But what I meant by "ask a chemist" is that a friend of mine who
is a chemist said that the vinegar and salt combined to form a
weak hydrochloric acid. I took his word for it.

--
Where ARE those Iraqi WMDs?

On Sat, 8 May 2004 06:02:31 -0700, "Paul O." <[email protected]>
wrote:

>Did a google for rust removal and saw a few references for removal of
>lightly rusted hand tools using table salt and vinegar. Any of you use this
>method? If this works would like to try it before scrounging up the parts
>for electrolytic rust removal. How much vinegar and salt do you mix with
>water? Thanks.

I've used sandpaper, then fine steel wool. Works just fine, and gets
rid of heavy to light rust.

Works on wood too, I hear.

On one very badly rusted item I started with a fairly decrepit rough
hone and scraped away with that. This was followed by sandpaper and
steel wool. It changed an old piece of rust I'd bought at a yard sale
for $5 Can into a cabinet maker's plane [hollow grooves in the base,
about 18" long.] I repainted also, and gave it as a gift to a person
with infinitely more talent than I have. He's still using it.

Bill.

On Sat, 08 May 2004 16:58:36 -0400, Bill Rogers <[email protected]>
wrote:

>a cabinet maker's plane [hollow grooves in the base,
>about 18" long.]

What's a "cabinet maker's plane" and why does it have a grooved base ?
--
Smert' spamionam

This may sound odd, but I have used plain ole brake fluid to remove rust. Wear
goggles when using it because it is death to eyes. Not really, but it is bad stuff.
No harm done to try it on some totally unimportant rusted steel or iron. "Nothing
ventured Nothing gained"
Hoyt W.

Ken Vaughn wrote:

> I have used the salt and vinegar method for cleaning small hand tools for
> several years. It works slowly, but works well for high carbon steel --
> especially plane irons and chisels. You will need to use something slightly
> abrasive to help clean the surface -- I have found that 3M ScotchBrite pads
> work well. On badly rusted plane irons, 2 to 3 cycles (soak for 30 minutes
> and scrub with pad) will clean it pretty well.
>
> The mixture is simple, standard 5% acidity white vinegar and table salt --
> just dissolve as much salt in the vinegar as it will take, and it takes
> quite a bit and dissolves slowly as it nears the saturation point. The
> mixture can be reused several times even though it turns red from the rust.
> Another important point, the steel will rust very quickly when removed from
> the solution. Keep some clear rinse water and a can of WD40 spray handy to
> clean and protect the tool.
>
> --
> Ken Vaughn
> Visit My Workshop: http://home.earthlink.net/~kvaughn65/
>
> "Paul O." <[email protected]> wrote in message
> news:[email protected]...
> > Did a google for rust removal and saw a few references for removal of
> > lightly rusted hand tools using table salt and vinegar. Any of you use
> this
> > method? If this works would like to try it before scrounging up the parts
> > for electrolytic rust removal. How much vinegar and salt do you mix with
> > water? Thanks.
> >
> > --
> > Paul O.
> > [email protected]
> >
> >

Aww gee,
Do you suppose it could be a #6 bailey?
I'd guess Stanley thought it should have grooves in the base (sole?).

"Andy Dingley" <[email protected]> wrote in message
news:[email protected]...
> On Sat, 08 May 2004 16:58:36 -0400, Bill Rogers <[email protected]>
> wrote:
>
> >a cabinet maker's plane [hollow grooves in the base,
> >about 18" long.]
>
> What's a "cabinet maker's plane" and why does it have a grooved base ?
> --
> Smert' spamionam

On Thu, 13 May 2004 09:04:26 -0700, Larry Blanchard
<[email protected]> posted:

>In article <[email protected]>,
>[email protected] says...
>> Ditto! All I can say is that the presence of ions in solution, ionic
>> strength, does definitely affect how species in solution react. Maybe there
>> is some physical chemistry website or ng you can visit and ask this
>> question. I'd be interested to know, too!
>>
>Wow! I didn't mean to start such a learned discussion :-). My
>knowledge of chemistry is limited to making various explosive
>compounds, learned long ago in my juvenile days. And lately, I
>think I've forgotten most of that - CRS seting in :-).
>
>But what I meant by "ask a chemist" is that a friend of mine who
>is a chemist said that the vinegar and salt combined to form a
>weak hydrochloric acid. I took his word for it.


Yep it makes what might be regarded as a "weak" hydrochloric acid.
This doesn't mean "dilute" BTW.
And guess what? Acetic acid is "weak" so there is effectively no
difference on this count. If you could make acetic acid "strong" you
would have the equivalent of hydrochloric acid.
Strong and Weak wrt acids means that they dissociate forming the H+
and anions to a greater or lesser extent.

On Thu, 13 May 2004 07:31:37 GMT, "Dan White"
<[email protected]> posted:

>"Sandy" <[email protected]> wrote in message
>news:[email protected]...
>
>> >>
>> >> >In article <[email protected]>, Sandy
>> >> ><> says...
>> >> >> I'm wondering why the salt. It surely does nothing but perhaps cause
>> >> >> problems later on.
>> >> >>
>> >> >See your friendly local chemistry professor :-).
>> >>
>> >> Do you not know what it does?
>> >> I've got a fairly reasonable handle on chemistry, and I can't see what
>> >> the salt is there for. The acetic acid will dissolve rust, and slowly
>> >> dissolve iron. No need for the salt unless you know of a good reason.
>> >> Please expound.
>> >>
>> >
>
>Sandy - I provided that rather long post that you responded to yesterday. I
>deleted it because of its size, and most of it is unrelated to your
>question. Also, you should realize (if you didn't) that I was passing on an
>article written by someone and didn't add any of my own thoughts, since I
>have forgotten so much chemistry I didn't want to dust off the brain cells.
>However, after reading your correct responses to his article, I had some
>thoughts that might possibly help, although I don't have a proven answer to
>your question. :(
>
>Let me say that I agree that the driving force in the reaction to dissolve
>the iron oxide is the hydronium ion, and formation of water on the right
>side of the equilibrium. I think the answer to why you have to add salt to
>the mix is that the salt increases the ionic strength of the solution. This
>improves the conductivity of the solution, and might have the effect of
>improving the rate of reaction by allowing an easier transfer of charges
>among the reacting ions. I know this is a bit vague, but I don't know if
>anybody has a really good answer on this.
>
>On the other hand, I recall that the equilibrium constant is really based on
>the activity of the ions in solution, and not their concentration. At
>higher ionic strengths (due to salt addition), I believe the activity, and
>therefore, the equilibrium constant, is lower than it would be without the
>salt. It seems to me this would tend to lower the dissociation of the
>acetic acid, and bind up even more H+ than without the salt. I think if
>this is correct, that it isn't the prevailing factor as I believe the salt
>does increase the reaction rate.
>
>I think the case of cleaning your copper pots is similar to the rust issue.
>In that case, if you pour vinegar on the copper pot with the oxides on it,
>you won't see much if anything happen. When you sprinkle salt on the
>surface wetted with vinegar, you quickly see the oxides disappear from the
>spots where the salt is dissolving. It is very clear that salt does speed
>the rate of reaction in the case of copper oxides, and I have to think it
>does the same with iron oxides.
>
>If you don't believe salt does anything, it is simple enough to test for
>yourself on two equal rusty spots. As to "why" it works, I have to think it
>is because of improved migration of electrons through solution due to the
>higher ionic strength of the solution. You might find, for example, that it
>takes two weeks to do the job with acetic acid alone, and one day with the
>salt added. Whatever problems you say the salt may cause later on may be
>outweighted by the time factor. Of course it is also possible (probable?)
>that the acid alone will not only take longer, but might in fact not remove
>as much oxide in the end.
>
>I had one other thought, improbable as it may be: I'd say the Fe3+ ion is
>relatively large compared to the Cl- ion. On the other hand, the acetate
>ion is of course a molecule and not a single ion. Maybe there is also a bit
>of steric hindrance going on as well. With no Cl- present, all the Fe3+ has
>to bond with the acetate molecule (3 of them). If there is some difficulty
>fitting 3 of these molecules on one Fe ion, the reaction could be inhibited.
>If free Cl- is present, it could more quickly and easily neutralize the Fe
>ion and move the molecule away from the reacting area.
>
>regards,
>dwhite
>

Interesting thoughts, Dan. This discussion has certainly disturbed
some old cobwebs in my attic :)
Gotta be over 40 years since that happened, although my daughter is
doing some elementary chem at night school and my helping her is also
raising the dust.

I don't disbelieve that it helps matters, but I just can't quite
understand why, although you have given me some food for thought.
I have never tested it, although as it happens, I had a rusty steel
band on a hose fiting that I wanted to remove and so stuck it in a
glass of white vinegar overnight. It cleaned the rust off enough to
show me that there was so much metal left that it needed the grinder
to get it off.

I can't really buy the conductivity argument, as what needs to happen
is the migration of H+ to the rust and the migration away from it of
iron ions.

I also can't buy the spatial problem of the size of acetate ions, coz
the stuff is all in solution, and not forming crystals.

BTW, the problems later with soaking a piece of rusty steel in salt
solution is that the salt gets trapped in the fine pits of the rusted
surface and later attracts moisture and THEN the conductivity of this
solution exposed to oxygen continues the corrosion at a pace.

Anyways, thanks for the stimulating discussion :)

"Dan White" <[email protected]> wrote in message
news:[email protected]...

> On the other hand, the acetate ion is of course a molecule and not a
single ion.

Minor correction: it is a single ion, just not a single atom.

dwhite

On Sat, 8 May 2004 06:02:31 -0700, "Paul O." <[email protected]>
wrote:

>Did a google for rust removal and saw a few references for removal of
>lightly rusted hand tools using table salt and vinegar.

Bad idea.

It doesn't work especially well.

It's _really_ risky on some copper-based alloys, especially bronzes.
You do _not_ want to start chloride-based corrosion on bronzes, as it
can be impossible to stop this afterwards.

Acid processes around ferrous materials are likely to give colour
changes (mainly darkening). You may find this useful, if you _want_ to
blacken things.

If you're cleaning ferrous metals, go for cathodic electrolysis. It's
easy to do, hard to get wrong, non-damaging and works really well.

For ferrous, a citric acid technique is commonplace (mainly for
museum-grade work on _really_ damaged items, particularly iron rather
than steel. IMHE, electrolysis is much simpler and works as well.

For brass or other cuprous alloys, a vinegar soak alone can be useful.
Don't leave it too long, or you'll find the brass de-zincs and turns
pinc.

--
Smert' spamionam

In article <[email protected]>,
Paul O. <[email protected]> wrote:
>Did a google for rust removal and saw a few references for removal of
>lightly rusted hand tools using table salt and vinegar. Any of you use this
>method? If this works would like to try it before scrounging up the parts
>for electrolytic rust removal. How much vinegar and salt do you mix with
>water? Thanks.

salt and any mild acid -- vinegar, lemon juice, etc. -- is extremely effective
at cleaning up oxidized copper. takes off the oxidization, without touching
the 'clean' metallic copper. I'm not sure of the entire chemical reaction,
but one of the byproducts is copper chloride, *bright* blue.

I'm guessing the situation with iron is similar, i.e.
iron oxide + salt ==> iron chloride + ??

I've used salt + lemon juice on copper cookware, for many years. lemon juice
straight out of the bottle, enough to wet the surface, and just sprinkle some
salt on. doesn't take a lot. The reaction is fairly quick -- seconds, to a
few tens of seconds. a single grain of salt seems to work over an area about
half the diameter of a penny. for heavy oxidation, a paper towel thoroughly
wetted with lemon juice, laid on the item, and then sprinkle some salt on.
The _nice_ thing about cleaning copper this way is that it is a 'self-limiting'
reaction. When the oxide is gone, things _stop_.


I've never played with the technique on iron,, but I'm guessing that iron
reacts considerably more slowly.

More than you probably wanted to know:

http://yarchive.net/metal/rust_remove.html

<snip>

"Sandy" <[email protected]> wrote in message
news:[email protected]...
> On Thu, 13 May 2004 08:21:20 GMT, "Dan White"
> <[email protected]> posted:
>
> >
> >"Dan White" <[email protected]> wrote in message
> >news:[email protected]...
> >
> >> On the other hand, the acetate ion is of course a molecule and not a
> >single ion.
> >
> >Minor correction: it is a single ion, just not a single atom.
> >
> >dwhite
>
> Yep, it's big, but this surely only matters when it has to associate
> with a cation. Crystalisation.

But steric hindrance is not a phenomenon reserved only to the crystalization
of compounds. There are several kinds of hindrance, and they don't all have
to do with crystallization. I don't think it really matters anyway because
the acetate ion probably isn't nearly large enough. It was just a thought.

dwhite

"Jim Martin" <[email protected]> wrote in message
news:[email protected]...

>>> What no one has discussed yet is the fact that it darkens the iron a
lot.<<<<

I would think that it might work something like bluing a gun. RM~

"Sandy" <[email protected]> wrote in message
news:[email protected]...
> On Thu, 13 May 2004 07:31:37 GMT, "Dan White"
> <[email protected]> posted:
>
> >"Sandy" <[email protected]> wrote in message
> >news:[email protected]...
> >
> >> >>
> >> >> >In article <[email protected]>, Sandy
> >> >> ><> says...
> >> >> >> I'm wondering why the salt. It surely does nothing but perhaps
cause
> >> >> >> problems later on.
> >> >> >>
> >> >> >See your friendly local chemistry professor :-).
> >> >>
> >> >> Do you not know what it does?
> >> >> I've got a fairly reasonable handle on chemistry, and I can't see
what
> >> >> the salt is there for. The acetic acid will dissolve rust, and
slowly
> >> >> dissolve iron. No need for the salt unless you know of a good
reason.
> >> >> Please expound.
> >> >>
> >> >
> >
> >Sandy - I provided that rather long post that you responded to yesterday.
I
> >deleted it because of its size, and most of it is unrelated to your
> >question. Also, you should realize (if you didn't) that I was passing on
an
> >article written by someone and didn't add any of my own thoughts, since I
> >have forgotten so much chemistry I didn't want to dust off the brain
cells.
> >However, after reading your correct responses to his article, I had some
> >thoughts that might possibly help, although I don't have a proven answer
to
> >your question. :(
> >
> >Let me say that I agree that the driving force in the reaction to
dissolve
> >the iron oxide is the hydronium ion, and formation of water on the right
> >side of the equilibrium. I think the answer to why you have to add salt
to
> >the mix is that the salt increases the ionic strength of the solution.
This
> >improves the conductivity of the solution, and might have the effect of
> >improving the rate of reaction by allowing an easier transfer of charges
> >among the reacting ions. I know this is a bit vague, but I don't know if
> >anybody has a really good answer on this.
> >
> >On the other hand, I recall that the equilibrium constant is really based
on
> >the activity of the ions in solution, and not their concentration. At
> >higher ionic strengths (due to salt addition), I believe the activity,
and
> >therefore, the equilibrium constant, is lower than it would be without
the
> >salt. It seems to me this would tend to lower the dissociation of the
> >acetic acid, and bind up even more H+ than without the salt. I think if
> >this is correct, that it isn't the prevailing factor as I believe the
salt
> >does increase the reaction rate.
> >
> >I think the case of cleaning your copper pots is similar to the rust
issue.
> >In that case, if you pour vinegar on the copper pot with the oxides on
it,
> >you won't see much if anything happen. When you sprinkle salt on the
> >surface wetted with vinegar, you quickly see the oxides disappear from
the
> >spots where the salt is dissolving. It is very clear that salt does
speed
> >the rate of reaction in the case of copper oxides, and I have to think it
> >does the same with iron oxides.
> >
> >If you don't believe salt does anything, it is simple enough to test for
> >yourself on two equal rusty spots. As to "why" it works, I have to think
it
> >is because of improved migration of electrons through solution due to the
> >higher ionic strength of the solution. You might find, for example, that
it
> >takes two weeks to do the job with acetic acid alone, and one day with
the
> >salt added. Whatever problems you say the salt may cause later on may be
> >outweighted by the time factor. Of course it is also possible
(probable?)
> >that the acid alone will not only take longer, but might in fact not
remove
> >as much oxide in the end.
> >
> >I had one other thought, improbable as it may be: I'd say the Fe3+ ion
is
> >relatively large compared to the Cl- ion. On the other hand, the acetate
> >ion is of course a molecule and not a single ion. Maybe there is also a
bit
> >of steric hindrance going on as well. With no Cl- present, all the Fe3+
has
> >to bond with the acetate molecule (3 of them). If there is some
difficulty
> >fitting 3 of these molecules on one Fe ion, the reaction could be
inhibited.
> >If free Cl- is present, it could more quickly and easily neutralize the
Fe
> >ion and move the molecule away from the reacting area.
> >
> >regards,
> >dwhite
> >
>
> Interesting thoughts, Dan. This discussion has certainly disturbed
> some old cobwebs in my attic :)
> Gotta be over 40 years since that happened, although my daughter is
> doing some elementary chem at night school and my helping her is also
> raising the dust.
>
> I don't disbelieve that it helps matters, but I just can't quite
> understand why, although you have given me some food for thought.
> I have never tested it, although as it happens, I had a rusty steel
> band on a hose fiting that I wanted to remove and so stuck it in a
> glass of white vinegar overnight. It cleaned the rust off enough to
> show me that there was so much metal left that it needed the grinder
> to get it off.
>
> I can't really buy the conductivity argument, as what needs to happen
> is the migration of H+ to the rust and the migration away from it of
> iron ions.
>
> I also can't buy the spatial problem of the size of acetate ions, coz
> the stuff is all in solution, and not forming crystals.
>
> BTW, the problems later with soaking a piece of rusty steel in salt
> solution is that the salt gets trapped in the fine pits of the rusted
> surface and later attracts moisture and THEN the conductivity of this
> solution exposed to oxygen continues the corrosion at a pace.
>
> Anyways, thanks for the stimulating discussion :)
>

Ditto! All I can say is that the presence of ions in solution, ionic
strength, does definitely affect how species in solution react. Maybe there
is some physical chemistry website or ng you can visit and ask this
question. I'd be interested to know, too!

dwhite

In article <[email protected]>, [email protected] says...
>
> I know nothing of chemistry so my questions are(1) When disposing of this
> stuff, anything in particular to worry about? (2) When storing, is a metal
> coffee can ok, should the container be vented? Don't want to use this
> mixture in anything that it will eventually eat thru. Thanks all for your
> help and advice.
>

I just store the stuff in the original vinegar container (labeled and in
the garage of course).

As for disposal, I just put the used stuff down the drain. The
ingredients are foodstuffs after all. The used mixture can stain from
the rust, so you do need to be careful that way.

"Larry Blanchard" <[email protected]> wrote in message
news:[email protected]...
>
> But what I meant by "ask a chemist" is that a friend of mine who
> is a chemist said that the vinegar and salt combined to form a
> weak hydrochloric acid. I took his word for it.
>

I think Sandy's point on this was that the H+ ions are already in solution
from the acetic acid. The presence of Cl- ions doesn't change anything
related to acid strength. For instance, it does not change the dissociation
constant for acetic acid, thereby increasing the H+ concentration. It is
the H+ concentration that determines how strong the acid is. My point was
that the activity of the ions in solution is decreased by the addition of a
lot of ions like Cl- and might actually lower the acidity. This does not
seem to be the predominant outcome, however, because the salt does increase
the speed of the reaction.

dwhite

"Paul O." <[email protected]> wrote in message
news:[email protected]...
>
> I know nothing of chemistry so my questions are(1) When disposing of this
> stuff, anything in particular to worry about? (2) When storing, is a metal
> coffee can ok, should the container be vented? Don't want to use this
> mixture in anything that it will eventually eat thru. Thanks all for your
> help and advice.
> --
> Paul O.
> [email protected]
>
>

I store mine in the 1 gallon plastic containers that the vinegar came in. I
don't think you would want to store it in a metal container. I haven't
disposed of mine yet, but I would think you could neutralize the weak acid
(vinegar) with baking soda. Coffee cans rust pretty quickly with just plain
water -- I use one with water for cooling lathe tools when I grind them --
it looks pretty rusty.

--
Ken Vaughn
Visit My Workshop: http://home.earthlink.net/~kvaughn65/

Hi Paul

As others have said, it works fine. What no one has discussed yet is the
fact that it darkens the iron a lot. Like to a dark brown color.

I don't really like that color for the surface of my jointer but I don't
dislike enough to have the tables ground.

Good luck,

Jim

"Paul O." <[email protected]> wrote in message
news:[email protected]...
> Did a google for rust removal and saw a few references for removal of
> lightly rusted hand tools using table salt and vinegar.

I know nothing of chemistry so my questions are(1) When disposing of this
stuff, anything in particular to worry about? (2) When storing, is a metal
coffee can ok, should the container be vented? Don't want to use this
mixture in anything that it will eventually eat thru. Thanks all for your
help and advice.
--
Paul O.
[email protected]

Sandy wrote:

> On Thu, 13 May 2004 22:15:37 -0500, Unknown <[email protected]>
> posted:
>
>>On Thu, 13 May 2004 09:04:26 -0700, Larry Blanchard
>><[email protected]> wrote:
>>
>>>,;In article <[email protected]>,
>>>,;[email protected] says...
>>>,;> Ditto! All I can say is that the presence of ions in solution, ionic
>>>,;> strength, does definitely affect how species in solution react.
>>>Maybe there ,;> is some physical chemistry website or ng you can visit
>>>and ask this
>>>,;> question. I'd be interested to know, too!
>>>,;>
>>>,;Wow! I didn't mean to start such a learned discussion :-). My
>>>,;knowledge of chemistry is limited to making various explosive
>>>,;compounds, learned long ago in my juvenile days. And lately, I
>>>,;think I've forgotten most of that - CRS seting in :-).
>>>,;
>>>,;But what I meant by "ask a chemist" is that a friend of mine who
>>>,;is a chemist said that the vinegar and salt combined to form a
>>>,;weak hydrochloric acid. I took his word for it.
>>
>>You shouldn't have as he was wrong.
>>
>>Vinegar is approximately 5% acetic acid plus some other goodies to
>>provide some taste. The hydrogen ion concentration is not sufficient
>>to react with metallic iron and therein lies one of the keys to the
>>process. The other key is the fact that chloride ions form a stable
>>complex with iron ions in solution. The iron chloride complex is
>>strong enough so that iron oxide will dissolve and form that complex.
>>Since there is no oxidant strong enough to react with iron metal the
>>net result is that the iron oxide goes into solution as the chloride
>>but the iron metal does not react.
>>
>>It is essential that the solution be kept oxygen free or the metal
>>will dissolve. This is particularly noticeable if you allow the metal
>>to be "derusted" to stick out of the solution into air e.g. you will
>>find that there has been a dissolution of iron metal at the air liquid
>>interface.
>>
>>The role of the acetic acid is to keep the solution acidic enough to
>>prevent the precipitation of iron oxide but low enough so that iron
>>metal does not react with hydrogen ions. It is the high concentration
>>of chloride that removes the rust not a "weak hydrochloric acid".
>
> How so exactly?
> Iron acetate is surely soluble enough?
>
>>If one used a concentrated salt solution without the acetic acid then
>>one would get a preciptate of hydrous iron oxide at the surface. This
>>would slow the reaction to a crawl.
>
> Um, surely without the hydrogen ions, you are not going to get any
> dissolution of anything in the first place. All you need is an anion
> along with the H+ that does not form an insoluble precipitate with the
> resulting iron ions.
>
>>A weak acid such as acetic acid allows one to put a lot of acid in the
>>solution but maintain a relatively low hydrogen concentration.
>
> Yep, that's what "weak" means wrt acids and bases.
>
>>The solution if kept covered can be used repeatedly until the amount
>>of dissolved iron reaches a point where the hydrous oxide begins to
>>precipitate.
>
> Where you have infact neutralised all the acetic acid present.
>
>>If the used solutions are left open to the air then it
>>will accumulate ferric chloride as a result of air oxidation.
>
> No, it will remain a solution of iron ions, acetate ions, sodium ions
> and chloride ions. The iron ions will slowly precipitate to iron
> hydroxide complexes as the final H+ ions are used up. No?
>
>>That
>>ferric chloride is an oxidizing agent strong enough to react with iron
>>metal which is the reason one gets an "etch line" at the liquid
>>surface.
>
> Yes, if what you had was ferric chloride. You don't. You have a
> neutral solution of the ions I just mentioned, surely.

You keep asking this question over and over. First, if you have ions they
are not "neutral". By definition an ion is electrically charged, hence it
is reactive. The solution is neutral because for each cation there is a
matching anion with the opposite charge, but the ions themselves are not
neutral at all. Second, chlorine is one of the most reactive of all
elements, hence any reaction involving chlorine will proceed at a higher
rate than one involving acetate. The end result is that by putting some
chlorine ions in the solution you end up with a faster reaction.

--
--John
Reply to jclarke at ae tee tee global dot net
(was jclarke at eye bee em dot net)

On Wed, 12 May 2004 13:43:42 -0700, Larry Blanchard
<[email protected]> posted:

>In article <[email protected]>, Sandy
><> says...
>> I'm wondering why the salt. It surely does nothing but perhaps cause
>> problems later on.
>>
>See your friendly local chemistry professor :-).

Do you not know what it does?
I've got a fairly reasonable handle on chemistry, and I can't see what
the salt is there for. The acetic acid will dissolve rust, and slowly
dissolve iron. No need for the salt unless you know of a good reason.
Please expound.

On Thu, 13 May 2004 22:15:37 -0500, Unknown <[email protected]>
posted:

>On Thu, 13 May 2004 09:04:26 -0700, Larry Blanchard
><[email protected]> wrote:
>
>>,;In article <[email protected]>,
>>,;[email protected] says...
>>,;> Ditto! All I can say is that the presence of ions in solution, ionic
>>,;> strength, does definitely affect how species in solution react. Maybe there
>>,;> is some physical chemistry website or ng you can visit and ask this
>>,;> question. I'd be interested to know, too!
>>,;>
>>,;Wow! I didn't mean to start such a learned discussion :-). My
>>,;knowledge of chemistry is limited to making various explosive
>>,;compounds, learned long ago in my juvenile days. And lately, I
>>,;think I've forgotten most of that - CRS seting in :-).
>>,;
>>,;But what I meant by "ask a chemist" is that a friend of mine who
>>,;is a chemist said that the vinegar and salt combined to form a
>>,;weak hydrochloric acid. I took his word for it.
>
>You shouldn't have as he was wrong.
>
>Vinegar is approximately 5% acetic acid plus some other goodies to
>provide some taste. The hydrogen ion concentration is not sufficient
>to react with metallic iron and therein lies one of the keys to the
>process. The other key is the fact that chloride ions form a stable
>complex with iron ions in solution. The iron chloride complex is
>strong enough so that iron oxide will dissolve and form that complex.
>Since there is no oxidant strong enough to react with iron metal the
>net result is that the iron oxide goes into solution as the chloride
>but the iron metal does not react.
>
>It is essential that the solution be kept oxygen free or the metal
>will dissolve. This is particularly noticeable if you allow the metal
>to be "derusted" to stick out of the solution into air e.g. you will
>find that there has been a dissolution of iron metal at the air liquid
>interface.
>
>The role of the acetic acid is to keep the solution acidic enough to
>prevent the precipitation of iron oxide but low enough so that iron
>metal does not react with hydrogen ions. It is the high concentration
>of chloride that removes the rust not a "weak hydrochloric acid".

How so exactly?
Iron acetate is surely soluble enough?

>If one used a concentrated salt solution without the acetic acid then
>one would get a preciptate of hydrous iron oxide at the surface. This
>would slow the reaction to a crawl.

Um, surely without the hydrogen ions, you are not going to get any
dissolution of anything in the first place. All you need is an anion
along with the H+ that does not form an insoluble precipitate with the
resulting iron ions.

>A weak acid such as acetic acid allows one to put a lot of acid in the
>solution but maintain a relatively low hydrogen concentration.

Yep, that's what "weak" means wrt acids and bases.

>The solution if kept covered can be used repeatedly until the amount
>of dissolved iron reaches a point where the hydrous oxide begins to
>precipitate.

Where you have infact neutralised all the acetic acid present.

>If the used solutions are left open to the air then it
>will accumulate ferric chloride as a result of air oxidation.

No, it will remain a solution of iron ions, acetate ions, sodium ions
and chloride ions. The iron ions will slowly precipitate to iron
hydroxide complexes as the final H+ ions are used up. No?

>That
>ferric chloride is an oxidizing agent strong enough to react with iron
>metal which is the reason one gets an "etch line" at the liquid
>surface.

Yes, if what you had was ferric chloride. You don't. You have a
neutral solution of the ions I just mentioned, surely.

"Unknown" <[email protected]> wrote in message
news:[email protected]...
> On Thu, 13 May 2004 09:04:26 -0700, Larry Blanchard
> <[email protected]> wrote:
>
> >,;In article <[email protected]>,
> >,;[email protected] says...
> >,;> Ditto! All I can say is that the presence of ions in solution, ionic
> >,;> strength, does definitely affect how species in solution react.
Maybe there
> >,;> is some physical chemistry website or ng you can visit and ask this
> >,;> question. I'd be interested to know, too!
> >,;>
> >,;Wow! I didn't mean to start such a learned discussion :-). My
> >,;knowledge of chemistry is limited to making various explosive
> >,;compounds, learned long ago in my juvenile days. And lately, I
> >,;think I've forgotten most of that - CRS seting in :-).
> >,;
> >,;But what I meant by "ask a chemist" is that a friend of mine who
> >,;is a chemist said that the vinegar and salt combined to form a
> >,;weak hydrochloric acid. I took his word for it.
>
> You shouldn't have as he was wrong.
>
> Vinegar is approximately 5% acetic acid plus some other goodies to
> provide some taste. The hydrogen ion concentration is not sufficient
> to react with metallic iron and therein lies one of the keys to the
> process. The other key is the fact that chloride ions form a stable
> complex with iron ions in solution. The iron chloride complex is
> strong enough so that iron oxide will dissolve and form that complex.
> Since there is no oxidant strong enough to react with iron metal the
> net result is that the iron oxide goes into solution as the chloride
> but the iron metal does not react.

Can you explain exactly what this iron-chloride complex is? Are you saying
that the iron oxide (rust) is preferentially breaking it's molecular bonds
and is reforming as some kind of complex, or as iron chloride? I take it
that it is not iron chloride because you say below that if oxygen is
introduced, then ferric chloride will form. Second question: What is the
reaction that transforms this "iron chloride complex" into ferric chloride?

Your mechanisms sound interesting, but it's hard to know if this is the
actual path without knowing the driving forces mathematically.

dwhite



>
> It is essential that the solution be kept oxygen free or the metal
> will dissolve. This is particularly noticeable if you allow the metal
> to be "derusted" to stick out of the solution into air e.g. you will
> find that there has been a dissolution of iron metal at the air liquid
> interface.
>
> The role of the acetic acid is to keep the solution acidic enough to
> prevent the precipitation of iron oxide but low enough so that iron
> metal does not react with hydrogen ions. It is the high concentration
> of chloride that removes the rust not a "weak hydrochloric acid".
>
> If one used a concentrated salt solution without the acetic acid then
> one would get a preciptate of hydrous iron oxide at the surface. This
> would slow the reaction to a crawl.
>
> A weak acid such as acetic acid allows one to put a lot of acid in the
> solution but maintain a relatively low hydrogen concentration.
>
> The solution if kept covered can be used repeatedly until the amount
> of dissolved iron reaches a point where the hydrous oxide begins to
> precipitate. If the used solutions are left open to the air then it
> will accumulate ferric chloride as a result of air oxidation. That
> ferric chloride is an oxidizing agent strong enough to react with iron
> metal which is the reason one gets an "etch line" at the liquid
> surface.
>
>

"Charles Erskine" <[email protected]> wrote in message
news:[email protected]...
> More than you probably wanted to know:
>
> http://yarchive.net/metal/rust_remove.html
>
> <snip>

This is the post I pasted in this thread originally to try and answer the
original question.

thanks,
dwhite

"Sandy" <[email protected]> wrote in message
news:[email protected]...
> On Wed, 12 May 2004 13:43:42 -0700, Larry Blanchard
> <[email protected]> posted:
>
> >In article <[email protected]>, Sandy
> ><> says...
> >> I'm wondering why the salt. It surely does nothing but perhaps cause
> >> problems later on.
> >>
> >See your friendly local chemistry professor :-).
>
> Do you not know what it does?
> I've got a fairly reasonable handle on chemistry, and I can't see what
> the salt is there for. The acetic acid will dissolve rust, and slowly
> dissolve iron. No need for the salt unless you know of a good reason.
> Please expound.
>

A quick search turned up this. This guy says it pretty well. The bottom
line is that the dissolution of the rust is prompted by the presence of
acetate and chloride ions. Iron would rather form a molecule with acetate
or chloride ions, which are soluble in water than remain as iron oxide,
which is insoluble. It is easy to introduce chloride in the form of salt,
and greatly increase the rate of reaction. The acid component of the
solution apparently does not take part in the dissolution of the rust (but
does dissolve the pure iron). It prevents the iron ions from reforming into
iron hydroxides, which are insoluble as is the original iron oxide (rust).

dwhite
From: [email protected] (Don Wilkins)
Newsgroups: rec.crafts.metalworking
Subject: Re: Rust
Date: Sat, 03 Apr 1999 01:35:37 GMT

On 2 Apr 1999 13:49:14 GMT, [email protected] (James W.
Swonger) wrote:

>In article <[email protected]>,
>Grant Erwin <[email protected]> wrote:
>>Muriatic acid works fine on rusty steel. It attacks the rust a lot
>>faster than
>>the steel.

As a chemist my first choice is hydrochloric acid as a solvent for
rust. Muriatic acid is readily available from the local hardware
stores. Yes it will also dissolve iron so you don't want to dump the
parts in and come back next week to see if the rust is dissolved.

>> ...
>>Vinegar alone is just a slower and milder form of the above. Commercial
>>pickling (submersion in a hot bath of a combination of strong acids) is
>>a more extreme version of the above.

I guess I want to mildly disagree with this. Vinegar is about 5%
acetic acid and also contains various organic stuff to provide some
taste but which are not particularly essential to the derusting.
Acetic acid is a weak acid and is going to behave differently than
hydrochloric acid.

The acetate ion forms soluble complexes with iron in solution which
aids in the dissolution. Since acetic acid is a weak acid there will
be a lot of the acetate ion tied up as undissociated acetic acid.

Hydrochloric acid is a strong acid and will be essentially completely
dissociated. The chloride ion forms a stable complex with trivalent
iron. It is the stability of this complex which aids in dissolving the
iron oxide (rust).

The addition of sodium chloride to the vinegar provides the extra
complexing power of the chloride to this mixture.


>>I'm not sure what happens chemically with salt and vinegar. I've only
>>used it to clean coppers, brasses and bronzes where if it works it is
>>like magic.

The chloride and/or the acetate ions form stable complexes with iron
ions which aids in the dissolution of rust. The trivalent iron forms
much more stable complexes which would lead one to believe that the
solutions would be more effective with e.g. Fe2O3 than with lower
oxides. Those tenacious black oxides are more dense and usually
contain divalent iron oxide.

> The combination gives you a mild hydrochloric acid with other junk
>floating in it.
>
>H+
>Na+
>Cl-
>CH3COO- (acetate)

Well sort of but there will also be a fair amount of undissociated
acetic acid.

>The acetate might give you some buffering effect, perhaps.
>I've seen acetic acid as an additive in a couple of electroplating
>recipes.

Yes but the buffering is not important in this case.

>The active principle is the HCl. For the copper based materials,
>it converts the insoluble oxides to soluble chlorides. I expect
>the same activity is responsible for iron cleaning.

In principle it is the same but it is the complexing power of the
anions (chloride and acetate) which drive the reaction. It needs to be
acidic enough to prevent the precipitation of the very insoluble iron
hydroxides. Other than that the hydrogen ion does not participate in
the dissolution of the rust. The hydrogen ion concentration does
contribute to the speed at which the metallic iron disappears.

There have been extensive discussions of an electrolytic method in the
past which I have stayed clear of. In that case it appears that the
rust removal is accomplished by creating hydrogen gas underneath the
rust coating and 'blasting" it off. This would explain why some of
those black dense oxide coatings are not removed.

The electrolytic method should not remove much metallic iron but on
the other hand what ever is on the surface as rust and gets "blasted"
off is not likely to be redeposited as metallic iron back on the piece
from the same location from which it originated.

In other words what is turned to rust ain't gonna get put back where
it once was.



Search for Google's copy of this article
----------------------------------------------------------------------------
----

From: [email protected] (Don Wilkins)
Newsgroups: rec.crafts.metalworking
Subject: Re: Rust
Date: Sun, 04 Apr 1999 18:25:07 GMT

On Sat, 3 Apr 1999 12:04:33 -0600, "Don Foreman"
<[email protected]> wrote:


>> >. It needs to be
>> > acidic enough to prevent the precipitation of the very insoluble iron
>> > hydroxides. Other than that the hydrogen ion does not participate in
>> > the dissolution of the rust. The hydrogen ion concentration does
>> > contribute to the speed at which the metallic iron disappears.
>
>Does this mean that using stronger acetic acid probably would not speed
rust
>removal but may increase rate of attack on steel?

I probably should say I don't know the answer to this question. My
opinion not backed by any research is that the 5% concentration of
acetic acid in vinegar probably isn't the optimum concentration. I ran
through some equations below to show why increasing the acetic acid
concentration doesn't increase the hydrogen ion concentration at the
same rate as occurs with hydrochloric acid. At 5% acetic acid I doubt
if the hydrogen ion concentration is high enough to dissolve iron.


Rust is already oxidized iron so when it is dissolved there is no
requirement for a change in oxidation state. When you dissolve the
metal (any metal) into an aqueous solution the metal is oxidized. If
the metal is oxidized then something must be reduced. If there is
nothing there to be reduced then the metal won't be oxidized.

In the mixtures under discussion (hydrochloric acid solutions or
vinegar-salt) there are three things present in the solution to serve
as oxidizing agents. These are the hydrogen ions from the acid, the
oxygen dissolved in the solution, and ferric iron.

So if you increase the hydrogen ion concentration then there will be
an increase in the dissolution of the metal (in this case iron). This
will be very apparent when a strong acid is used such as hydrochloric
acid because all of the acid is ionized.

Acetic acid however is a weak acid so an increase in the concentration
of this acid does not increase the hydrogen ion concentration
proportionally. If you look at the equation for the ionization of
acetic acid

HAc <----> (H+) + (Ac-) one finds that there is an equilibrium
constant for that dissociation

(H+) (Ac-)/ (HAc) = ~10^-5

since (H+) = (Ac-)

(H+) = square root of (HAc)*10^-5

With that it is obvious that increasing the acetic acid concentration
is not going to cause a rapid increase in the dissolution of iron
metal.

Who ever came up with the addition of chloride ion did so to provide
something to form a complex with the dissolved iron among other
things. If you look at that equation above for the dissociation it is
obvious that if they had added acetate ion (e.g. sodium acetate) it
would have reduced the hydrogen ion concentration probably causing
precipitation of iron compounds. Clever idea for kitchen chemicals.

On the other hand there are other reactions which will come into play
once the process is started.

As most of you know ferric chloride can be quite corrosive and is a
pretty good etching agent. The ferric chloride produced from the
dissolved rust is going to look at that freshly produced iron metal
surface as a nice place to do some etching (dissolve iron metal).

Of course there will be some divalent iron present as well. This will
come either from the rust or by the reaction of the ferric chloride
with the metallic iron.

Oxygen (from air) dissolved in the solution will oxidize the ferrous
ions up to ferric. This ferric can then react again with metallic
iron. If you don't want to dissolve iron metal then you should remove
air from the procedure. A little bit of ferric chloride in a sodium
chloride solution with a continuous supply of oxygen can dissolve one
hell of a lot of iron.

I hope this hasn't confused the issue more than it helped. I suspect
that one could make a better rust removal solution using glacial
acetic acid, water, and sodium chloride. I believe the formula was
developed because of the ready availability of vinegar and salt. I
would vary the chloride concentration as well as the acetic acid
concentration. If you don't want to dissolve iron metal keep oxygen
out of the process.

If you are inclined to do this be careful with glacial acetic acid. It
doesn't burn or cause immediate discomfort on the skin but if not
removed promptly will cause large thick patches of skin to come off
leaving tender exposed meat. This is a type of chemical burn that you
are not likely to get more than once.

As I read and reread this post I am not entirely satisfied but there
is some very complex chemistry taking place and it is not easy to
describe it in a short note.