Making Calcium Oxide & Calcium Hydroxide
WARNING: This procedure
involves high heat and open flame, as well as production of a caustic
choose to attempt this or any of the experiments or procedures
this site, you do so entirely at your
oxide (CaO) and calcium hydroxide (Ca(OH)2) are related
compounds that have great usefulness in chemistry.
CaO and Ca(OH)2 are available from chemical supply houses,
it's interesting to explore their preparation from raw materials.
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
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
bacterial urease, which starts converting urea to ammonia and making
the odor worse (see http://www.crscientific.com/experiment2.html);
CaO, however, will kill bacteria and denature enzymes with
which it comes in direct contact.
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.
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.
goggles MUST be worn at all times when doing the procedure or
the compounds, even when just opening their containers. Gloves
are also a good idea.
find and read the MSDS for calcium oxide; JT Baker has one at this link.
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
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.
in particular mustn't be stored in moist areas or near acids or
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
into the lungs. The same dust precaution goes for Ca(OH)2.
Calcium carbonate (CaCO3)
Water (pref. distilled and de-gassed)
Beaker made of heat-resistant glass
Crucible w/ lid
Ringstand and clay triangle
Acetylene torch, Charcoal-fired kiln, or Bunsen burner
Methods & Observations:
those safety goggles on at all times!
a polycarbonate face
shield and a pair of welder's gloves while heating the crucible and its
I. Preparation &
Roasting of Limestone
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
obtain. Black, brown, or dark gray limestone contains more
organic matter, iron compounds, and other
using chalk instead, it should be
pure calcium carbonate; some varieties contain magnesium
carbonate, etc.; these are not suitable. Although calcium
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
our starting material.
carbonate should be in powdered form. If not already powdered, it
be pulverized in a
steel mortar and pestle or crushed to dust with a hammer. Safety
goggles must be worn
during the whole experiment.
rock particles have a peculiar habit of going straight for the eyes.
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.
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
a liter-sized container's worth could take a day or two of constant
to drive off all the CO2.
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
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.
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.
a small, charcoal-fired
kiln made of firebricks and refractory mortar is good
for making larger amounts of quicklime. The kiln's "lid" consists
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
When preparing only a few grams of CaO at a time, a crucible and bunsen
burner are of course easier.
crucible lid should be either ajar or off completely
during the first 30 seconds or so of heating, in order that any
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
it there for about five minutes.
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
said before, both compounds must be stored where children and other
unqualified people cannot get to them.
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.
make calcium hydroxide on purpose, the CaO is "slaked" with water, slowly and in the
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
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
get to it.
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-
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.
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
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
and nail clippings could instead be placed in a spot plate well or a
beaker and let stand in a humid atmosphere at 35-40°C; this
would also release ammonia, which can then be detected by a moistened
strip of pH paper, litmus paper, or phenolphthalein paper.
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
oxide reacts violently with acids, even weak ones, when water is
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
• Calcium oxalate from Ca(OH)2 and oxalic acid.
again, the reader is cautioned never to add unslaked lime (CaO)
acids. It must be cautiously hydrated into Ca(OH)2
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
is somewhat caustic and
must be treated with care. Do NOT get it in the eyes.
In order to use this website or any information
be advised that malicious or negligent use of caustic materials may
invite criminal penalties.
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Amateur Scientists has
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Merck Index, 10th
Edition. Rahway, New Jersey: Merck and Company, Inc., 1983.
CR Scientific Catalog