This is the old Caveman Chemistry website.

Please visit the new website at

A Primitive Alkali: Potash


Fire has been quite useful to us, both for the production of warmth and for the conversion of materials from one form to another (e.g. clay to pottery). Even the ashes left over from the fire turn out to have useful properties. Early people discovered that wood ashes could be used for cleaning. This may seem counterintuitive to us, since we usually associate ashes with dirt. But wood ashes are not simply dirt. The major components of wood ashes are potassium carbonate (potash) and sodium carbonate (soda ash). From a chemical standpoint these two compounds are very similar. So similar that while ashes have been used for millennia the difference between sodium and potassium carbonate was only recognized in the 19th century. The elements in the first column of the Peridic Table, containing sodium and potassium, are called the alkali metals.

Potash was man's first base. A base is a material which removes hydrogen ions (H+) from aqueous solution. We also refer to such a material as an alkali. Alkali's have a bitter taste an a slighltly soapy feel when rubbed in the fingers. The isolation of alkali is the precursor to making true soap.

Even today potassium and sodium carbonate remain extremely important chemicals. U.S. Production of sodium carbonate alone was 9 billion kg making it the 11th most-produced chemical in the U.S.

The Chemistry of Potash

You will recall from the metathesis project that all sodium and potassium compounds are soluble in water. When an ionic compound dissolves in water it dissociates into cations (positive ions) and anions (negative ions). For potash we write:

(I) K2CO3(s) -----> 2 K+(aq) + CO32-(aq)

Before we proceed we need to discuss a little of the chemistry of water itself. Like ionic compounds, water has the capacity to ionize, that is, to break apart into ions. In any sample of water a tiny fraction of the water molecules will be ionized according to the reaction:

(II) H2O(l) <-----> H+(aq) + OH-(aq)

The double arrow signifies that water is in equilibrium with hydrogen ion and hydroxide ion. At any given time there will be lots of water, a little hydrogen ion, and a little hydroxide ion present. In pure water at room temperature, for example, 1 water molecule out of 556 million will be ionized. Because one water splits into one H+ and one OH-, the H+ and OH- concentrations are equal in a neutral solution. When an ionic compound dissolves in water, it may absorb or release hyrogen or hydroxide ions. If the net effect is that hydrogen ions outnumber hydroxide ions, the solution is called acidic. If there are more hydroxide ions, than hydrogen ions the solution is called basic or alkaline.

We have already seen our first acid. When wine or mead "goes sour," bacteria oxidize ethanol to acetic acid. Acetic acid ionizes to produce hydrogen ions and acetate ions:
CH3COOH(aq) <-----> H+(aq) + CH3COO-(aq)
The net effect is that hydrogen ions are produced and hence the solution is acidic. An acidic solution tastes sour.

When posassium carbonate dissolves in water, it ionizes into potassium and carbonate ions, as we have seen. Once in solution, the potassium ions simply float around and are very unreactive. We call such an ion a spectator ion because it does not participate in further chemical reactions. The carbonate ion, on the other hand, acts as a base in aqueous solution, that is, it absorbs hydrogen ions wherever it finds them. The reaction we write depends on whether the solution was acidic or basic to begin with. In acidic solution we write:

(IIIa) CO32-(aq) + H+(aq) <-----> HCO3-(aq)

while in alkaline or neutral solution we write:

(IIIb) CO32-(aq) + H2O(l) <-----> HCO3-(aq) + OH-(aq)

The carbonate ion will react with whatever species is around. If there is lots of hydrogen ion (acidic solution) the hydrogen ion sticks to the carbonate ion and forms a new ion, the hydrogen carbonate ion, or bicarbonate ion. If there is not much hydrogen ion around, carbonate in effect steals a hydrogen ion from water, leaving a hydroxide ion behind and producing an alkaline solution. Notice that once again we have used a double headed arrow to denote an equilibrium rather than a reaction that goes to completion.

There are two more reactions that further complicate the picture. First, if the solution is really acidic, the bicarbonate ion can also stick to a hydrogen ion:

(IVa) HCO3-(aq) + H+(aq) <-----> H2CO3(aq)
(IVb) HCO3-(aq) + H2O <-----> H2CO3(aq) + OH-(aq)

H2CO3 is called carbonic acid. Carbonic acid can can decompose into water and carbon dioxide gas:

(V) H2CO3(aq) <-----> H2O(l) + CO2(g)

You may have noticed this phenomenon in fermenting mead. The yeast produces carbon dioxide gas. You can detect the production of this gas because when the bottle is sealed, it becomes pressurized. Only a thin layer of bubbles appears on top of the mead. But if the cap is opended, the mead blossoms into a foamy brew teaming with bubbles. Before your opened the cap, the carbon dioxide gas was dissolved in the mead. When the cap was opened and the pressure released, carbon dioxide came out of solution, escaping as a gas.

This sequence of equilibria is also active in carbonated beverages. When water is pressurized under carbon dioxide gas, some of the gas dissolves. Some of the dissolved gas combines with water to form carbonic acid. Some of the carbonic acid splits into hydrogen ion and bicarbonate ion. Some of the bicarbonate ion further splits into hydrogen ion and carbonate ion. All five species will be present but depending on the acidity of the solution there will be more of one than another.


It is useful to be able to quantify acidity. Scientists use the pH scale to do this. pH (pee Aych) is defined as the negative of the common logarithm of the hydrogen ion concentration. In mathematics, the common log is used to count the order of magnitude of a number, i.e. to count the number of zeros. For example, in pure water, the H+ concentration is 0.0000001 moles/L. There are seven zeros in this number, and the pH is 7.

The flip side of this specification is the OH- concentration and the negative of the common log of this concentration is called pOH (pee Oh Aych). Because water can ionize into H+ and OH-, it turns out that pH + pOH = 14. In an acidic solution there are lots of H+ ions floating around but not very many OH- ions, and the pH is low. In a neutral solution, there are equal numbers of H+ and OH- ions and the pH is 7. And in an alkaline solution, there are few H+ ions, lots of OH- ions, and the pH is high. Of course, the concentration of the ions depends on the concentration of the solution. So to make a fair comparison, we should compare solutions with the same total concentration. Here are the pH and pOH values for several solutions with a concentration of 1% (i.e. 1 gram of compound in 100 grams of solution):

CompoundDominant SpeciesH+ ConcentrationpHOH- ConcentrationpOHTest Paper
Acetic acidCH3COOH0.000140.000000000110
Carbonic acidH2CO30.0000150.0000000019
Pure WaterH2O0.000000170.00000017
Potassium BicarbonateHCO3-0.0000000180.0000016
Potassium CarbonateCO32-0.0000000001100.00014

One of the most convenient ways to measure pH is with an indicator. An indicator is a substance which changes color when the pH changes. There are many indicators. You may have noticed, for example, that tea turns from dark brown to tan when lemon juice (an acid) is added. We will use test paper which changes from blue (alkaline) to green to yellow (neutral) to orange to red (acidic).

Recrystalization as a Purification Technique

One of the most fundamental problems in chemistry is that of purifying a substance. We have seen that most things in nature are actually mixtures, either homogeneous or heterogeneous. But to have any control over chemical reactions, a chemist must first be assured that his starting materials are pure. Later in the course we will look at distillation as a technique for purifying substances that are gases or liquids under ordinary conditions. For solids, however, we turn to recrystalization as our primary purification technique.

We have seen that a mineral, particularly a crystalline mineral, is essentially a pure substance, either element or compound. How are such minerals formed in nature? One possibility is that the mineral crystals cooled from molten rock. Another is that the crystals formed from substances dissolved in water. It is this mechanism which we will try to exploit in this project.

In the metathesis project we discussed solubility as if it were a yes or no proposition. In reality it is a little more complicated. You know from everyday experience that sugar is soluble in water. If you add a single grain of sugar to 2 liters of water, the sugar will all dissolve--no solid sugar will remain. And if you add a second grain, it too will dissolve. In fact, if you add a teaspoon of sugar, the entire amount will dissolve. You may add a second and a third teaspoon, but anyone who has ever added sugar to tea or coffee knows that there comes a point where the water is saturated with sugar, that is, all the sugar that can dissolve, has dissolved. Any sugar in excess of this amount will just settle to the bottom as a solid. The amount of a material which will dissolve in a given amount of water is called its solubility. Here are solubilities of some common substances:
SubstanceSolubility (grams/100 mL water)
potassium carbonate147331
calcium chloride75159
sodium chloride3639
potassium chloride3457
sodium carbonate22421
calcium carbonate0.0010.002

Suppose, then, that I start with an ancient sea. Dissolved in the seawater are all kinds of things dissolved out of the soils and rocks in the area which drains into the sea. Some of these substances, like sodium chloride, have high solubility while others, like calcium carbonate, have low solubility. Eventually, the geological conditions change, the sea is cut off from the ocean, the drainage patterns change, and the sea begins to dry out, perhaps over the course of hundreds of thousands of years. The Dead Sea and the Great Salt Lake are two modern examples of such a situation. Now, as the sea evaporates it becomes more concentrated until it becomes saturated in the least soluble material it contains. Like the excess sugar added to tea, this least soluble material falls to the bottom and is deposited as a layer, perhaps calcium carbonate. As the evaporation continues, the substances present are deposited in reverse order to their solubilibies. Finally, the most soluble substances present are deposited as the uppermost stratum as the sea gives up the last of its moisture. The substances deposited depend on what was present in the original sea, and the order in which they are deposited depends on their relative solubilities.

One more hitch in the story is that solubility depends on temperature. Notice that while the solubility of sodium chloride is about the same in hot and cold water, the solubility of sodium and potassium carbonate is much greater in hot water than in cold water. We will exploit this property in separating the soluble carbonates from the other components of wood ash.

Wood Ashes

Whatever we extract from wood ashes must be there to begin with. Wood ashes are a complex heterogeneous mixture of all the non-flammable, non-volatile minerals which remain after the wood and charcoal have burned away. Because of the presence of carbon dioxide in the fire gases, many of these minerals will have been converted to carbonates. Burned soil may also be present. So the ashes probably contain predominately sodium and potassium carbonate, sodium and potassium chloride, silica, and calcium carbonate.

If we add the ashes to water, the soluble potassium and sodium salts will dissolve while the insoluble silica and calcium carbonate will settle to the bottom. We can then drain off the water (containing the "good stuff") and throw the insoluble material away. To separate the chlorides from the soluble carbonates, we will exploit the greater solubility of the carbonates in hot water. We will bring the liquid to a boil and continue boiling until enough water boils away for an insoluble precipitate to form. This is very likely a mixture of sodium and potassium chloride. From this point, we will continue boiling until half of the remaining water is removed. At this point we can be reasonably certain that only the soluble carbonates remain in solution. We will carefully pour off the hot liquid into another container, leaving the solid material behind. As the liquid cools to room temperature, the less soluble sodium carbonate will precipitate leaving the more soluble potassium carbonate in solution. Finally, the remaining solution can be drained off and boiled to dryness, producing solid potassium carbonate.

One of the observations you make should be that it takes a lot of wood to make a little ash and a lot of ash to make a little potash. Thus, while it is not particularly difficult to extract potash from wood, you will go through an enormous amount of wood to produce commercial amounts (pounds and tons) of potash. This will have implications for us later in the semester.

Other Potash Pages

Potash Quiz

The potash quiz consist of three questions on any of the following topics discussed in this page.

Safety and Common Sense

Just because potash and soda ash are "natural" and you are extracting them from wood ashes, doesn't mean they are safe. Just because they are chemicals with chemical names and formulas doesn't make them dangerous. When thinking about chemical hazards, you should always consider the amount and concentration of the substance in question. Potash makes up only a small percentage of wood ashes, which are not particularly hazardous. But as the potash is extracted, concentrated, and purified, it becomes more deserving of care.

Potash and soda ash are relatively strong alkali's. The are moderately caustic, which means they will damage skin. Consequently, you should not rub them all over your body or get them in your eyes and you should not eat them. Nevertheless, they don't warrant paranoia. If you get some on your skin, wash it off. You should wear glasses to protect your eyes, but if some gets in, you should splash cold water into your eyes. You should not eat it. It's not good for you. But a little taste will do you no harm. If you are foolhardy enough to eat several teaspoons of it, call the local poison control center (VA 1-800-451-1428).

Information on chemical hazards is summarized in a Material Safety Data Sheet for each compound. These sheets often tell you more than you want to know, but they are worth looking at.


Our goal is to extract as much of the soluble carbonates from wood ashes as possible while leaving behind the insoluble components. You will need a couple of handsful of wood ashes, some water, our old friend, the 2 L soft drink bottle, and a pottery bowl, non-aluminum saucepan, or glass beaker. The first container (2 L bottle) need only hold water. We use the soft drink bottle for convenience. The second container needs to be fire and water proof. If we were going for historical accuracy, we would use pottery. But any pan or beaker that can be used on a stove will work. Do not use an aluminum pot. Aluminum reacts with strong bases and your project will be ruined.

Place your wood ashes into the 2 L bottle until it is about 1/3 full of ashes. Fill the rest of the bottle with hot water, place the cap on the bottle and shake it up. The soluble carbonates (as well as any other soluble materials) will dissolve while the insoluble silicates, carbonates, aluminosilicates, and any other insoluble materials will settle to the bottom. Any charcoal present will float to the top. Place the bottle where it will not be disturbed and let it sit overnight.

The next day you should find that the sediment has settled to the bottom, the charcoal is floating on the top, and the water in between is clean and clear. Remove the cap and give the bottle a gentle squeeze to force the charcoal up and out of the bottle. Then carefully pour about 1 L of the clear water into a pan or beaker. Stop pouring before any of the stirred-up sediment reaches the mouth of the bottle. What you should have now is a pan with about 1 L of what looks like clear water. If you taste it and it tastes a little bitter or soapy, you are on the right track.

Place the pan onto the stove or hotplate and place a spoon or glass rod into it to prevent superheating (which can lead to splattering). Bring the water to a boil and continue boiling until all the water has evaporated. A little bit of scale or fine grey powder will remain in the pan. Let it cool and then scrape it into a container. This is your product. It contains all of the soluble materials which were present in the the ashes to begin with. This could include sodium and potassium chlorides, sulfates, hydroxides, and carbonates. Of these, only the hydroxides and carbonates are basic. I will test your product with pH test paper to determine whether it is alkaline.

If you were interested in further purifying your product, you could recrystalize it again. This time you would start with your crude product (instead of ashes), dissolve it in water, boil it until it was almost dry, and filter it while hot to remove any materials less soluble than the carbonates. You would then allow it to cool and the carbonates would precipitate out leaving anything more soluble still in solution. By repeated application of this procedure, you could even separate sodium carbonate from potassium carbonate. But for our purposes, you crude potash should be alkaline enough. To test it for yourself, just taste it. It should taste bitter, like soap.

Criteria for Success

You will have already passed the potash quiz when you bring your potash for evaluation. Your potash must be grey or white, with no obvious contamination. A wet pH test paper should turn blue when touched to your potash. Of course, if you fail, you can try again (once per day) until you pass.

Sodium carbonate is sold as washing soda. Other synonyms are soda ash and soda. You can buy washing soda in some grocery stores, on the laundry detergent aisle, next to the borax. It is used to remove calcium from hard water by forming an insoluble precipitate of calcium carbonate.

Sodium bicarbonate, a much milder alkali than sodium carbonate, is sold as baking soda. It is also called sodium hydrogen carbonate. It is used in baking, as the name suggests, and is a common ingredient in antacid tablets. You can buy baking soda in the grocery store on the baking goods aisle. In the presence of acids, it gives off carbon dioxide, which makes bread and cake nice and fluffy. Baking powder contains baking soda along with an acid.