From: firstname.lastname@example.org (Casey Donovan)
Subject: Re: Old NaClOH solutions still good?
Date: 13 Jan 1999
On Tue, 12 Jan 1999 17:56:53 -0800, summers@remove_alum.mit.edu (David
P. Summers) wrote:
>Does anyone know of commerical solutions of sodium hypochlorite
>in water (ie bleach) keep their efficaciousness over long periods
>of time (years)? (ie is a bottle that is years old "still good"?)
They don't. Any bottle of household bleach, typically 5.25% wt. NaOCl
(not NaClOH), over a year old is just alkaline salt water. Industrial
concentrations, 10.0% to 12.5% or so, last a little longer in cool
temperatures, but not dramatically so.
Charts of decomposition vs. temperature aren't hard to find, but a
rough rule is that household bleach should be used within 2 - 4 months
at air conditioned home temperatures. A corollary of this rule -- buy
the cheap store brand. It's made to the same initial concentration as
the advertised stuff (except for the very pricey "Ultra" type) and it
usually is fresher.
From: email@example.com (Casey Donovan)
Subject: Re: available chlorine - what is it?
Date: 02 Aug 1996
firstname.lastname@example.org (Steven Westwood) wrote:
>Can anyone give me an explanation of what "available chlorine" refers
>I have a bottle of sodium hypochlorite which is stated to be
>5% available chlorine. How does this equate with the concentration
>of sodium hypochlorite?
>Thanks in advance for any answers
Available chlorine is a "trade term" for the concentration of sodium
hypochlorite or hypochlorite ion in the solution, expressed as Cl2.
Since Cl2 has a molecular weight of 70.91, and sodium hypochlorite a
formula weight of 74.45, 5 % available chlorine means
5. x (74.45/70.91) = 5.25 % NaOCl.
Why is it expressed as Cl2 and not just Cl? Because it yields a bigger
number! :-)) The justification is that molecular (gas) chlorine can thus
be considered to be (have) 100 % available chlorine. When gas chlorine is
dissolved in water, it hydrolzes to form one equivalent of hypochlorous
acid plus one equivalent of hydrochloric acid:
Cl2 + H20 ------> HOCl + HCl.
Only the hypochlorous acid is "available" as an oxidant or disinfectant.
The hydrochloric acid half just reduces the pH. When the solution is
neutralized with sodium hydroxide, a mixture of sodium hypochlorite and
sodium chloride is produced, but still only the sodium hypochlorite is
useful as a disinfectant.
From: email@example.com (Casey Donovan)
Subject: Re: available chlorine
Date: 03 Aug 1996
>5% available chlorine should mean 50 g/L of Cl2.
>some chlorine is not available because it is chloride.
This statement is true, but could easily be misunderstood.
It would be better to say 50 g/L (really 50 g/Kg) of Cl, expressed as Cl2.
If 50 g of chlorine is dissolved in 1 kilogram of water solution, 36.99 g
of HOCl is formed (plus 25.70 g of HCl). Since the chlorine content of
HOCl is 67.6 %, only (36.99 x .676) or 25 g of Cl is available as an
oxidant. That looks like, and is, 2.5 % of the solution. However, since
the trade term "available chlorine" expresses Cl in terms of Cl2, the
solution is said to have 5 % available chlorine.
It's a little like expressing the Mg content of water in terms of CaCO3.
100 mg/L of Mg is often stated as 412 mg/L, as CaCO3.
>BTW: how about free and total chlorien in tap water?
Free chlorine in water is Cl present as hypochlorite ion, OCl-.
Combined chlorine in water is Cl "combined" into a reaction product with
some reductant such as ammonia or another amino compound. In this form, it
has greatly reduced capacity as an oxidant or disinfectant. Chloramines
are the primary culprits in "chlorine" taste, odor, and eye irritation.
Total chlorine is the sum of Free and Combined.
Subject: Re: bleach question
Date: Wed, 11 Aug 1999 14:11:26 GMT
In article <firstname.lastname@example.org>,
"Dane Myers" <email@example.com> wrote:
> how is it that bleach takes the color out of things?
Colored substances fall into two categories:
organics and transition metals. Bleaches don't work
with most transition metal stains (e.g. rust)
but only with organic stains.
Most organic chemicals are colorless.
Visible light covers the wavelength range 400-700 nm
approximately. Compounds with only single bonds
in them only absorb in the deep ultraviolet near 200 nm
or in the infrared beyond 1500 nm. Compounds with
double bonds in them absorb at longer wavelengths
but still not above 300 nm. To get long wavelength
absorption you need large systems of conjugated
single and double bonds. However even this does not
guarantee color, e.g. benzene and naphthalene,
which are both a colorless. To get an organic
molecule which absorbs in the visible is quite
Systems of conjugated single and double bonds
do tend to be fairly stable, but to get strong
absorption in the visible you also need heteroatoms,
ususally oxygen and nitrogen, and these are points
of weakness at which a bleach can attack.
The active ingredient in ordinary bleach is
hypochlorite, OCl-, which is a good oxidizing
agent. Chances are when hypochlorite is through
oxidizing a molecule, it doesn't absorb light
any more. Of course there are always exceptions,
as Barry Hunt has pointed out.
Most bleaches are oxidizing agents.
The modern ones tend to be peroxide-type, since
they don't smell. Sulfite or
bisulfite is also used as a bleach, in
paper manufacture I think,
and it's a reducing agent.
Subject: Re: bleach question
Date: Mon, 16 Aug 1999 08:31:25 GMT
In article <1999Aug13.firstname.lastname@example.org>,
email@example.com (Matthew Sugdon) wrote:
> Why is it unusual for an organic molecule to absorb
> in the visible. Take a look at all those pale yellow
> solutions sitting around your lab, n-pi*
> transitions, sure it's not bright but it's still colored.
> What about dyes? Aren't they organic molecules? They
> are colored. So for you to say that organic molecules
> aren't colored is way off the mark, the Colour
> index lists some 10,000 organic colored molecules
> with thousands more have been synthesised but are
> unlisted Benzene and napthalene are colorless but
> what about the more complex aromatics? They show
> some pretty funky deep blues when you conjugate 9 or 10
I don't think you've really contradicted anything I said.
There are at least a million known organics, and I
think you'll grant me that most dyes are formed
from conjugated systems with oxygen or nitrogen
either in the system itself or close to it. So from
your own number, 1% of organics are strongly colored.
The whole organic industry started with Perkin's
synthesis of mauveine, about 200 km southwest
of you (www.ch.ic.ac.uk/motm/perkin.html). There
is considerable financial incentive to synthesize
dyes, which have higher added value than most
other organics, and I think that this skews
the 1% ratio of dyes:other organics a bit.
> Bleaches in general decolorise dyes by oxidation,
> azo dyes for example can be cleaved quite nicely
> at the azo bond.
My statement was that the weak points
where a bleach can attack are the heteroatoms.
I can't see where you said anything different.
> just look at your clothes after
> a few washes with a bleach-containing detergent,
> typically these sorts of bleaches are sodium
> perborate, sodium percarbonate, hydrogen peroxide.
Je n'ai jamais prétendu le contraire.
> When most people think of bleach they think of
> something really powerful like hypochlorite,
> this will take the colour out of most dyes, however
> some polycyclic vat dyes are stable so could be used.
> You then have the problem of coloured effluent, how
> do you decolourise such a stable molecule? what is so
> appealing about a colored bleach? All you are doing
> is throwing money down the toilet!
I don't disagree with you there! But nobody
listens to ascetics in modern marketing departments.
And if you're rich, why not light your cigars
with $100 bills?
My solution would be to get the color in this
case from a transition metal. I always think
permanganate looks so pretty, a pity its stability in
aqueous solution leaves something to be desired.
From: firstname.lastname@example.org (John Vinson)
Subject: Re: NaClO + H2O2 -> ??
Date: Wed, 10 Jun 1998 20:33:48 GMT
On Wed, 10 Jun 1998 17:24:17 GMT, "BRIAN L. DOSS"
>I am a first year chem student and was just wondering exactly how this
>simple reaction proceeded. The two components (NaClO & H2O2) are common
>household chemicals. NaClO is bleach and H2O2 is hydrogen peroxide. Does
>this reaction produce Cl gas or H gas? In all honesty, I only want to
>know this because I am curious. If anyone answers this ?, Ithank you in
Neither; it produces oxygen by oxidizing the peroxide to O2:
H2O2 + OCl- --> O2 (g) + Cl- + H2O
At least a portion of the oxygen produced is in the excited singlet
state (ground state oxygen is a triplet); this phosphoresces as it
converts to the ground state, giving off a deep red glow. Do it in
the dark with your eyes dark adapted - one way that works is to use a
buret with hypochlorite, dripping into unstirred 3% peroxide; you get
readily visible red flashes.
Use proper caution with both chemicals, of course: if you try the
above do it in the light first, and don't break the buret trying to
grab it in the dark!