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Xioa
March 15th, 2004, 02:46 AM
I searched if there was already a topic on this, and I was surprised to discover there wasn't. This post deals with using red cabbage as a cheap and easily accessible pH indicator. This method is pretty widely known but has not been discussed almost at all within this forum, even though it seems to be very useful.
The pH indicator is prepared by blending red cabbage in a blender, straining off the liquid, making sure there is no plant matter in it, then put it into your solution you are testing the pH of. The red cabbage indicates at a pH range of 1 to 12. Here's a picture of the color range, from rhodium's site:
http://www.rhodium.ws/chemistry/equipment/pictures/ph-cabbage.jpg
I hope this helps some of you that are having trouble finding a good pH indicator.

frogfot
March 15th, 2004, 06:41 AM
Whow, that seems to be an accurate scale (except for pH 3-7), this is a very useful thing.
Any idea what is the active substance that changes color? Maby one could extract it and concentrate, to make it storageable since cabbage juice doesn't store well.

mongo blongo
March 15th, 2004, 02:59 PM
IIRC you can make it with rose petals too.

Wild Catmage
March 15th, 2004, 03:50 PM
http://home.howstuffworks.com/question439.htm - This site claims that the colour change is caused by one pigment contained in the red cabbage.

This site claims that the colour change is caused by two different pigments - http://www.penpages.psu.edu/penpages_reference/29503/29503200.HTML

For information on extracting anthocyanins, go to - http://www-saps.plantsci.cam.ac.uk/osmoweb/chloro1.htm

Tuatara
March 15th, 2004, 05:58 PM
I've done quite a lot of messing with natural pH indicators (mostly to entertain the kids) and I've found practically any intensely coloured flower or berry will yield a useable indicator. Even red wine works! Doubtless different sources will produce indicators for different pH ranges, sadly I have no calibrated pH meter to test them with.

The biggest drawback with red cabbage is the smell :D

Xioa
March 15th, 2004, 06:50 PM
I've read that the red cabbage juice can be stored for 2 or 3 weeks if you refrigerate it. Heating the juice destroys the pigment, so if you plan on keeping it a while, dont put it in the sun, keep it cool.
Also, Tuatara, you do not have to use a pH meter to measure the pH, just use a strong acid or base (weak ones are useable too, if you get the Ka/b of them) and then record the moles you add until it turns color, then calculate the pH at which it changes color. It requires a bit of extra work though :rolleyes:

Tuatara
March 15th, 2004, 09:11 PM
Heating didn't destroy my red cabbage - thats how I made the extract! Just boiled the stuff for 15mins (yech! what a stink). Boiling also kills off any bacteria, so the resulting liquid should keep really well if covered and refrigerated.

T_Pyro
March 16th, 2004, 04:24 AM
Anthocyanin is the water-soluble pigment that lends a red or purple colour to fruits. Since the active component of the indicator is the one that is coloured, it follows that any fruit or flower having the same pigment will also act as an indicator. This (http://www.agsci.ubc.ca/courses/fnh/410/colour/3_22.htm) page has some more information regarding the effects of change in Ph on Anthocyanins.

frogfot
March 17th, 2004, 06:08 AM
It showed to be very easy to extract pigment from red beet. I've chopped one beet in food processor and then extracted two times by boiling it with water. On addition of big ammounts of alcohol (cheap denaturated), red goo precipitated (possibly with rest of inorganic salts). After filtering, solvent was slightly orange. Dryed product is noncrystalline dark red solid (<1 g), which looks like somekind of rosin.
Now, I prepared a concentrated solution of this red solid, gonna see how well it stores..
Also, time to by some red cabbage ;)

vulture
March 17th, 2004, 09:53 AM
Store it as a diluted solution, indicators are already effective in very low concentrations.

T_Pyro
March 17th, 2004, 03:08 PM
I've been experimenting with black grapes (purple=anthocyanin most probably), and found that the indicator turns green in the presence of a base, and dark red, turning orange in the presence of an acid.

Frogfot, is the "red goo" that precipitated a concentrated form of the pigment? I didn't understand how the substance precipitated out... The red dye from beetroot is water soluble, as are the possible inorganic salts, so what was it that actually precipitated?

EDIT: Looking at the Ph chart above, it seems that the Ph range of the black grape would be more like 1-2=last 2 test tubes on the scale 3-14= first 12 test tubes. How this shift in colour with Ph takes place, is a mystery!

frogfot
March 17th, 2004, 03:40 PM
I said it precipitated on addition of alcohol.. since anthocyanin is composed mainly of benzene rings, I assumed it would have bad solubility in cold alcohol. And I assumed that nearly all pigment was precipitated since filtered solution became orange..
One might think that this pigment have different color in alcohol.. anyway, I tested to dilute the orange filtrate and the red color didn't come back, so I must've precipitated it all.

As I understand from link you've posted before, pH indicator should contain a big conjugated double bond system, also conjugated with H+ donors/acceptors (like -OH and =O respectively). So it seems that it should be pretty easy to produce own indicators.. we actually made one when some of us tryed to produce phorone, the byproducts that was synthesized fall in the above-discribed creteria (with big conjugated doublebond system), this stuff was responcible in color change when acidic reaction mix was neutrolised.
Anyway, there should be other simple reactions that can produce similar complicated compounds..

T_Pyro
March 17th, 2004, 04:47 PM
I tried doing the same thing with the black grape extract, using isopropyl alcohol (Don't have any ethyl aclohol), but the solution just got diluted a bit. How much alcohol did you have to add till the precipitation began?

The conjugated system with carbonyl carbons and hydroxy groups help to stabilise the carbocation by resonance (delocalisation of electrons). Ph indicators are generally buffer solutions using alkali metal salts of weak acids, or salts of strong acids and weak bases which have a characteristic colour. Eg. In Mythyl orange:
http://img34.photobucket.com/albums/v104/T_Pyro/Methyl-Orange.gif
In the presence of a stronger acid, the -SO3- gets converted to -SO3H.

frogfot
March 18th, 2004, 06:27 AM
To this boiled down extract I added approx 10 volumes alcohol (maby more), try cooling to -20*C. There are possibility that grapes have other pigments..

T_Pyro
March 18th, 2004, 04:44 PM
A little OT, but... I don't think the product produced in the Phorone synthesis would fall into the category of "indicators". More likely, it was just a condensation product like:
http://img34.photobucket.com/albums/v104/T_Pyro/Reaction-scheme.gif
And so on, producing a polymeric chain.

The final product wouldn't be a salt of a strong acid+ weak base, or strong base+ weak acid, hence it can't act as an indicator. However, I agree with what you said about producing other artificial indicators. Since an alkali metal salt of a coloured organic weak acid would act as an indicator, I was thinking maybe it could be possible to create other indicators by coupling 2 or more benzene rings with a diazo bond (to provide the colour to the indicator), with a sulfonyl group on one of the rings to act as the weak acid. Or else, the weak acid part could be a phenolic -OH bonded to a ring having strongly deactivating groups. Also, the compound would have to be soluble in water to act as an acid/base indicator.

vulture
March 18th, 2004, 05:58 PM
I've probably mentioned this before, but if you're titrating something with either NH3 or HCl, CuSO4 can be used as an indicator for neutral pH.

The smallest excess of NH3 will cause the solution to go to blue, a small excess of HCl will cause it to turn green. I've tried this and the transition is rather sharp and easily detected by the naked eye.

A small amount of coppersulfate will do.

frogfot
March 19th, 2004, 11:17 AM
Thats a useful thing, with copper sulfate :)

T Pyro, it doesn't have to be a salt. The compound changes color because its double bonds change position. And this can happen simply if a carbonyl group takes an H+ (in acidic medium), or as you said, hydroxyl group loses it's H+ in strong basic conditions. This affinity to H+ can vary greatly depending on how large is the conjugated system.
IIRC, absorbed light depends on total energy of doublebond system.

Oh, and ketone (or aldehyde) condensation mechanism is a bit different from that pic, because, in third step, the H+ will leave methyl group to form a double bond much faster than that carbonyl oxygen would attack.

T_Pyro
March 19th, 2004, 10:09 PM
Using copper sulfate, shouldn't the end point be indicated by a pale blue ppt. due to the formation of copper hydroxide? The dark blue colouration appears only after there's enough excess of ammonia for the entire copper to be precipitated out, and then dissolved as a complex. By that point, the ammonia would be in excess, beyond the equivalence point. Nickel chloride, however, might work better because with ammonium hydroxide, it readily forms a blue complex, even before forming any ppt. The result is a sharp green to dark blue transition.

Frogfot, what you're talking about seems a lot similar to the basis of infrared spectroscopy. So if we have a compound such as 2,4,6 unsaturated ketone (or aldehyde), would it act as an indicator? It does sound a bit doubtful to me... :(

About the condensation mechanism, it's quite different from the condensation due to the presence of bases. However, if the H+ ion was extracted from the methyl group, due to keto-enol tautomerism, it would speedily revert to acetone. On the other hand, in the presence of a concentrated, strong acid, the concentration of the carbocation species in the third stage just might be high enough to promote the reaction.