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CR Scientific

Making Calcium Oxide & Calcium Hydroxide

WARNING: This procedure involves high heat and open flame, as well as production of a caustic material.  If you choose to attempt this or any of the experiments or procedures described on this site, you do so entirely at your own risk.


Calcium oxide (CaO) and calcium hydroxide (Ca(OH)2) are related compounds that have great usefulness in chemistry.
While CaO and Ca(OH)2 are available from chemical supply houses, it's interesting to explore their preparation from raw materials.
Calcium oxide is commonly known as "quicklime", "burnt lime", "unslaked lime", or sometimes just "lime".  While CaO is sometimes available on the consumer market,  it is getting harder to find.  Most stores now sell the less-useful calcium carbonate (CaCO3), which is powdered limestone, not lime;  specifically asking for "quicklime" at farm stores nowadays will often get you puzzled looks.  Since limestone cannot denature proteins in the manner of quicklime, it is generally useless as a stall disinfectant and deodorizer, for example.  CaCO3 can make it just alkaline enough to activate bacterial urease, which starts converting urea to ammonia and making the odor worse (see; CaO, however, will kill bacteria and denature enzymes with which it comes in direct contact.
CaO is representative of group II oxides (MgO, SrO, etc).  It can be prepared by intense roasting of the corresponding carbonate or hydroxide;  the formation is reversible in the presence of CO2 or H2O, respectively.  Make sure to read the safety section about CaO and water.
Ca(OH)2, also called "hydrated lime" or "slaked lime", is somewhat soluble in water and is useful not only as a laboratory reagent, but also for making several other useful compounds.  It isn't as strong a protein denaturant as CaO, but it must still be treated with respect.


Safety goggles MUST be worn at all times when doing the procedure or handling the compounds, even when just opening their containers.  Gloves are also a good idea.
Please find and read the MSDS for calcium oxide;  JT Baker has one at this link.
Hydrating CaO with water (known as "slaking") generates a great deal of heat and converts the compound into Ca(OH)2.  Slaking is potentially dangerous because it can cause spattering, burns, glass breakage, or even unwanted combustion of nearby organic materials.  Don't mix CaO with acid;  the reaction is violent and will spatter hot acid all over you.
Although neither compound is toxic in the same sense as chromium, lead, nickel, and so forth, CaO and Ca(OH)2 are strong bases and can therefore still cause death, injury, or blindness if handled carelessly.  They can denature proteins, dissolve cell membranes, disrupt bodily pH, and generally cause burns;  therefore, they must be stored in a securely-locked area where children can't get to them. 
CaO in particular mustn't be stored in moist areas or near acids or water.  It must be handled as carefully as lye, with added precaution not to disturb the fine CaO powder up into the air where it could contact the eyes or find its way into the lungs.  The same dust precaution goes for Ca(OH)2.


Calcium carbonate (CaCO3)
Safety goggles
Welding gloves
Water (pref. distilled and de-gassed)
Beaker made of heat-resistant glass
Crucible w/ lid
Ringstand and clay triangle
Crucible tongs
Acetylene torch, Charcoal-fired kiln, or Bunsen burner

Methods & Observations:
Keep those safety goggles on at all times!
Wear a polycarbonate face shield and a pair of welder's gloves while heating the crucible and its contents.

I.  Preparation & Roasting of Limestone

The first step is to acquire and prepare a suitable quantity of calcium carbonate.  If using limestone as the source, the stone should be the whitest shade that one can obtain.  Black, brown, or dark gray limestone contains more organic matter, iron compounds, and other contaminants. 
If using chalk instead, it should be pure calcium carbonate;  some varieties contain magnesium carbonate, etc.;  these are not suitable.  Although calcium sulfate is an ingredient in some chalks and will decompose at very high temperatures to give CaO and SO2 (or SO3 under some conditions), it's easier and safer to use pure calcium carbonate as our starting material.
The carbonate should be in powdered form.  If not already powdered, it has to be pulverized in a steel mortar and pestle or crushed to dust with a hammer.   Safety goggles must be worn during the whole experiment.  Small rock particles have a peculiar habit of going straight for the eyes.
We'll need enough crushed CaCO3 to fill a porcelain crucible about 1/2 to 3/4 of the way.  The amount is not critical as long as it the vessel isn't full;  the powder level should be at least a centimeter below the rim of the crucible.
Calcium carbonate decomposes at 825 C to yield carbon dioxide and calcium oxide.   The amount of time required to keep it at this temperature depends on how much material is present.  A small crucible's volume of powder may require only a few minutes of roasting, while a liter-sized container's worth could take a day or two of constant heating to drive off all the CO2
An air-acetylene torch or a charcoal-fired kiln can provide enough heat to decompose calcium carbonate.  An air-acetylene torch can easily heat a 15-mL porcelain crucible on a clay triangle to about 850-900 C, hot enough to melt sterling silver and therefore hot enough to roast CaCO3 into CaO.  An oxy-acetylene torch is too hot, though;  localized heating from such a flame could destroy the crucible.   A propane torch, on the other hand, may not be quite hot enough when we consider the heatsink effects of the crucible and its support ring. 

Roasting calcium carbonate in a mini kiln
Above:  a small, charcoal-fired kiln made of firebricks and refractory mortar is good for making larger amounts of quicklime.  The kiln's "lid" consists of a few bricks placed on top to slow the escape of heat.  Note the steel pipe blowing air into the kiln through a hole in the bottom.   The air source is the blower on an old vacuum cleaner.  Internal temperature of this kiln is more than enough to decompose calcium carbonate into CaO.  Unlike a torch flame, the kiln holds in the heat and has a higher "practical" temperature.  There is of course the danger of unseen moisture or air pockets in the mortar, which could pose a danger when it's fired.  The refractory mortar should cure for at least a month before firing.

When preparing only a few grams of CaO at a time, a crucible and bunsen burner are of course easier.

The heating must be done in an outdoor area or on a fireproof lab bench with adequate ventilation;  in both cases there must not be any distractions, clutter, or combustible materials nearby.  No children should be present or anywhere nearby.
The crucible lid should be either ajar or off completely during the first 30 seconds or so of heating, in order that any moisture will be driven off.  After that, put the crucible lid on with a pair of tongs and keep the lid on during the rest of the heating.  The lid must fit loosely enough that expanding gases can still escape.  Heat the crucible to redness and keep it there for about five minutes.
Do not attempt to cool the crucible with anything;  let it cool by itself in air.   Keep the lid on while the crucible is cooling.   Take the lid off only when necessary;  the influx of moisture- and CO2-laden air should be kept to a minimum.

II.  Storing Calcium Oxide; Making & Storing Calcium Hydroxide
As said before, both compounds must be stored where children and other unqualified people cannot get to them. 
Calcium oxide will react with water, producing calcium hydroxide and a great deal of heat.  If you want it to remain calcium oxide for very long, it must be put in an airtight container where moisture and carbon dioxide can't get in.  Otherwise it will slowly turn to calcium hydroxide and calcium carbonate on standing.   If you have a desiccator, store the CaO in there.

Above:  it's helpful to store CaO in a desiccator.   Just make sure it's
somewhere it can't be knocked over.  Never draw vacuum on the
desiccator when there are closed containers inside.

To make calcium hydroxide on purpose, the CaO is "slaked" with water, slowly and in the following manner:
 CaO is added, a little at a time, into an excess of water (Do not add the water into the CaO !).  The heat generated from hydration of CaO can sometimes crack glass;  in extreme cases it can cause paper to catch fire.  Use a borosilicate beaker to do the mixing.  (Keep those safety goggles on; remember that CaO can permanently damage the eyes.)  When it's done slaking, allow the water to evaporate in your locked laboratory where no people or animals can get to it.
To prepare the solution of Ca(OH)2 known as limewater, simply don't evaporate the liquid that results from the previous step.  Let any insoluble matter settle to the bottom;  decant and save the liquid.  The pH of a saturated Ca(OH)2 solution is 12.4 at room temperature (Merck Index, 1983), having a concentration of about 0.01 M.   Calcium hydroxide is only sparingly soluble, but it is a strong base (i.e., it ionizes almost completely into Ca++ and OH- ions).
Ca(OH)2 is not as reactive as CaO to moisture and air, but it can still absorb CO2 and slowly form the carbonate.  It must therefore be kept tightly sealed.


The preparation of CaO and Ca(OH)2 are conceptually very simple, as we've seen.  Naturally, some readers will be asking what these compounds are good for.  Let's examine some uses.

Calcium oxide reacts with ammonium salts or with proteins to liberate ammonia gas (NH3) .  An experiment from a 1960's Chemcraft set had the experimenter place fingernail clippings and calcium oxide together in the hand.  CaO against bare skin would be considered poor form nowadays;  even though the caustic calcium oxide has difficulty penetrating the outer, dead skin of the hand, it's better to avoid skin contact altogether.  The CaO and nail clippings could instead be placed in a spot plate well or a micro beaker and let stand in a humid atmosphere at 35-40C;  this would also release ammonia, which can then be detected by a moistened strip of pH paper, litmus paper, or phenolphthalein paper.
Since calcium oxide has such a strong affinity for atmospheric moisture, it can be used as a desiccant (calcium chloride is preferable for this, however).  It also acts as a CO2 absorber, especially when mixed with granulated sodium hydroxide (lye).  This mixture is known as soda-lime and is often used in special columns through which gases are fed in order to remove the CO2 and water vapor they might contain.  Soda-lime is also important for some organic chemistry procedures, such as the traditional preparation of methane (CH4) from sodium acetate.  It's also used for removing the -COOH group from aromatic carboxylic acids such as benzoic and salicylic.
Calcium oxide reacts violently with acids, even weak ones, when water is present.

Calcium hydroxide can react to produce other useful compounds.  For example:
• Calcium chloride from Ca(OH)2 and dilute hydrochloric acid;
• Calcium nitrate from Ca(OH)2 and dilute nitric acid;
• Potassium hydroxide from Ca(OH)2 and potassium carbonate;
• Calcium acetate from Ca(OH)2 and acetic acid; 
      Sodium acetate from calcium acetate and sodium carbonate.
• Calcium oxalate from Ca(OH)2 and oxalic acid.

Once again, the reader is cautioned never to add unslaked lime (CaO) directly to acids.  It must be cautiously hydrated into Ca(OH)2 first.
The solution of Ca(OH)2, known as limewater, is also used to detect carbon dioxide.  If a gas thought to contain CO2 is bubbled into the liquid, a cloudy, white precipitate of calcium carbonate means there is CO2 (the unknown gas mixture of course shouldn't contain anything that reacts violently with water...) 
Limewater is somewhat caustic and must be treated with care.   Do NOT get it in the eyes.

Notice:  In order to use this website or any information contained herein, you must read and agree to the Terms of UsePlease be advised that malicious or negligent use of caustic materials may invite criminal penalties.

Copyright:  The Society for Amateur Scientists has permission to reprint this article in full.  The article otherwise remains copyright of CR Scientific and may not be copied, reproduced, mirrored, or distributed without prior written permission (click here for contact info).  If you find any of the information on this page useful in your own research, please include the appropriate citation(s) in your work. 

Work Cited:

Merck Index, 10th Edition.  Rahway, New Jersey: Merck and Company, Inc., 1983.

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