GEOS 110 Lab 7
Precipitation of Calcium Carbonate & Salts
Names_______________
Calcium carbonate, CaCO3 is the major chemical component of many minerals and rocks. These
include limestone, travertine, dolostone, marble, calcite, aragonite, coral, some marine shells,
stalactites, stalagmites. Precipitation of CaCO3 forms these substances. Dissolution of CaCO3
allows the raw materials to be transported in solution to be re-precipitated elsewhere.
(CaCO3) precipitates from solutions that contain calcium ions (Ca2+) from the chemical
weathering and dissolution of minerals like feldspar, pyroxene, amphibole or calcite itself and
carbonate ions (CO32-) from CO2 dissolved in rain water as carbonic acid, or bicarbonate ions
(HCO3)- . The carbonate ion can enter the solution in the form of carbon dioxide (CO2) from the
air that is dissolved in the water:
CO2 (g)
+ H2O (l) → H2CO3 (aq)
↔
CO32- (aq)
+ 2H+(aq) (1)
Carbon dioxide + Water forms Carbonic acid which dissociates to Carbonate ion + Protons
Carbonic acid is a weak but abundant natural acid. That makes most rain slightly acidic so that it
falls with a pH of ~5.6 where 7.0 is neutral. Other acids occur naturally including: hydrochloric,
sulfuric and nitric acid from volcanic gases, acid generating minerals like pyrite, and organic
acids from decomposition and oxidation of plant matter. Acid falling or flowing over rocks and
minerals is the first part of the rock cycle for chemical weathering. Because of this weak and
most abundant carbonic acid, hydrogen ions are available to substitute for cations as they
chemically attach themselves to mineral surfaces. The acidic protons make portions of the
minerals more soluble. Most minerals react with acid slowly and dissolve to make soluble ions
(aqueous salts) and clays or oxide/hydroxide residues (soils). Even Quartz SiO2 can dissolve this
way as follows:
SiO2 (s) + 2H2O ↔ H4SiO4 (aq) or
SiO2 (s) + 4 H+ +4 e- ↔ H4SiO4 (aq) (2)
Quartz + Water react to Tetrahydrosilicic Acid or Quartz “Hydrolizes to” ….
When natural waters or their solutions are cold and acidic, there is a lot of dissolved molecular
carbonic acid and bicarbonate ions from equation (1). As waters heat up, dissolved gas is less
soluble and CO2 gas comes out of solution. When waters are alkaline, the equilibrium shifts
further to the right and there are lots of dissolved carbonate ions.
This carbonate ion can then react with calcium ions that are already in the water:
Ca2+ (aq) + CO32- (aq) ↔ CaCO3 (s)
(3)
Other common ions like Na+, K+ or Mg+2 are too soluble to precipitate until nearly complete
evaporation. But calcite will form in your pipes, hot water heater or even in your kidneys! In
nature, most carbonate forms in hot, tropical marine shelves (reefs, limestone), in hot springs
(travertine), carbonate cements (in sedimentary rocks or speleothem cave formations) or in
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evaporating salty shallow seas or in playa lakes (salt deposits from salars/salinas). This serves to
take carbonate ions (which after all are soluble CO2 gas) and tuck them away as minerals and
rocks. If they are buried beneath warm salty water or other sediments beneath the tropic seas and
kept away from acid rain, (or cold acidic sea water in the deepest ocean) they can stay there and
serve to reduce the environmental load of CO2 in the air. This kind of natural carbon storage like
the growth of the Great Barrier Reef off Australia, will eventually reduce the CO2 in the air but,
left to its own slow pace, this will takes tens of millions of years just by weathering of rocks and
precipitation of limestones alone. Meanwhile we are making more by burning more fossil fuels
and slash in tropical rain forests. We also heat up crushed limestone to get lime (CaO), Ca+2 for
Portland cement and release even more CO2 this way! This latter industrial activity from burning
fossil fuels and decrepitating limestone generates 5% of the CO2 produced globally.
A. Experiment
Equipment: calcium chloride (CaCl2 solution), sodium carbonate (Na2CO3) solution; graduated
cylinders, beakers, ring stand, filter paper, aluminum foil, binocular microscopes, natural
calcium carbonate; weak hydrochloric acid in bottles; sea water; streak plates, evaporate
minerals; Bunsen burners
Procedure: Measure out 10 mL of aqueous calcium chloride, CaCl2 solution in a graduated
cylinder. This provides the Ca2+ ions as in equation (3) above. In a separate graduated cylinder,
measure out 10 mL of sodium carbonate, Na2CO3 solution. This provides the CO32- ions as in
eqn (3) above.
1. Observe the nature of the separate solutions and describe them here. ____________________
______________________________________________________________________________
___________________________________________________________________________ (2)
2. Write the total ionic equation for the 2 solutions you brought together and any products you
produced. Note: Both the Cl- and Na+ are highly soluble and will be present both before and
after the reactions. They are “spectator ions”, just there for the charge balance and along for the
ride! Write your equation here:
___________________________________________________________________________ (5)
3. Pour both solutions into the same small beaker. Record your observations. ________________
_____________________________________________________________________________
__________________________________________________________________________ (2)
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The solid is a precipitate of calcium carbonate, CaCO3. This compound has low solubility in
water and readily precipitates out of the solution. Although the reaction is not exactly the same
as most precipitations in nature, it demonstrates the nature of solid forming from liquid, without
any freezing. CaCl2 (aq) + Na2CO3 (aq)  CaCO3 (s) + 2NaCl (aq) (4)
Similar reaction can occur between briny formation water and fresher runoff, hot geyser fluids
and fresher groundwater or between runoff and playa lakes.
4. It is possible to separate the calcium carbonate from the sodium chloride remaining in solution
in the water by filtration. Write your name on the edge of a piece of filter paper with water fast
ink. Weigh it and record the weight of your paper to 2 decimal places. Filter paper ______ g (1)
Procedure: Fold a filter paper in half, then in half again. Open it up as a conical shape and fit it
into a funnel. Support the funnel in a small ring (second drawer in bench) attached to a lab stand
(middle cupboard in bench).
a. Pour the mixture through the filter, and collect the liquid in a beaker. We will use this next so
save it! The residue in the filter paper is the calcium carbonate. Remove the filter paper and open
it out and lay it face up with your crystal residue on a paper towel to wick away any excess salt
water. Spread the filter paper on a piece of aluminum foil and to allow the CaCO3 to dry on a hot
plate for a few minutes at low heat. Weight your paper dry with CaCO3 ______________g (1)
b. While your filter paper and precipitate are drying out, evaporate your little beaker or
evaporating dish of <20 mL of solution. According towhat remains after reaction 4 above, this
should be just salt water. Once it dries out and cools off, examine and describe the crystal residue
under a hand lens or binocular microscope and draw and describe the crystal shapes you see.
There are samples of all the common evaporate minerals in the lab for you to compare to.
(2)
Draw them here:
5. Obtain a separate sample of natural calcium carbonate, powder it on a streak plate and test its
reaction to a few drops of weak hydrochloric acid. Chemically, this will add protons back onto
the carbonate ion, making some carbonic acid, which in turn will then dissociate back to CO2 and
water as follows:
CaCO3 (s) + 2 HCl (aq) → H2CO3 (aq) + CaCl2 (aq)
(5)
The carbonic acid then does equation (1) backwards:
H2CO3 (aq) → CO2 (g) + H2O (l)
(1R)
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a. Compare a pinch of your precipitate to natural calcite powder, side by side on a dark streak
plate, and describe the results. This should convince you that you have precipitated Calcite which
is the first mineral salt to form on evaporation of any hard briny warm natural waters.
Compare the reaction of the natural calcite powder to your precipitate: ____________________
_____________________________________________________________________________
___________________________________________________________________________ (2)
6. Compared to the inorganic precipitation of Calcite we performed, look at representative
natural limestones and note some of their more complex structures, textures and the nature of
their grains. Examine and describe each rock briefly in the table below.
(20)
Rock Name
Grain
Composition
Structures
Textures
Biology
Fossils
Other
Characteristics
1. Coquina
2. Fossiliferous
Limestone
3. Dolostone
4. Chalk
5. Travertine
7. Which of the rocks above represent simple inorganic precipitation of calcite as in our mixed or
evaporated solutions and comment on how common these types are? ______________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
______________________________________________________________________________
___________________________________________________________________________ (2)
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8. Reflect on the other rocks from this set and comment on other biogeochemical cycles and
discuss their significance to a.) the rock cycle, b.) the carbon cycle). What sort of systems and
system interactions (hydrosphere, atmosphere, biosphere, geosphere) do these rocks represent
and how long it might take to form them (or destroy them)? ____________________________
____________________________________________________________________________
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9. The Seas are all salty with about 35 per mil salinity. In the 1880’s, Usiglio did an experiment
taking a bucket of Mediterranean sea water and he evaporated it to dryness. Solubility and
precipitation are empirical phenomena (WYSIWYG). This is because water as a strong dipole
arranges itself around the ions for various salts in particular geometries that change with:
salinity, temperature and the composition of other ions in solution. Precipitation is basically a
dehydration reaction, which requires the removal of solvation water from around and between
ions before they can collide and precipitate as a solid. Strangely enough, NaCl is not the first
mineral to form. The order Usiglio found was:
Usiglio’s experiment
1. CaCO3
Calcite
2. CaSO4 or CaSO4 – 2H2O
Anhydrite or Gypsum
3. NaCl
Halite
4. MgSO4
Epsomite (Kieserite)
5. MgCl
Bischofite
6. Complex K-Mg-Cl-SO4-OH & KCl Bitterns: Carnallite, Langbeinite, Polyhalite,
Kainite, respectively: (KMgCl3 * 6H2O , K2Mg2(SO4)3 , K2Ca2Mg(SO4)6 * H2O ,
KMg(SO4)Cl * 3H2O)
7. KCl
Sylvite
In addition to these minerals, Usiglio found small amounts of iron oxy-hydroxides and NaBr
which more usually forms impurities as does NaI, in Halite. Waters on land are far more varied
in their chemistry due to the composition of the local bedrock and minerals undergoing
weathering. Complex carbonate chlorides, sulphate chlorides and borates are often present as
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well as salts of rarer elements like Lithium, Strontium, Barium etc. These constitute only traces
in marine waters whose chemistry is dominated by the submarine weathering of basalt.
a. It is geologically common to find the 1st 3 minerals from Usiglio’s list in evaporate deposits
but exceedingly rare to find anything in the lower part of the list. For this reason those salts are
the most valuable. Examine the evaporate minerals from our collection. Some of them are quite
delicate, so try not to break or dissolve them. Find minerals from the list above and describe how
they are different.
(28)
Salt formula
Mineral Name
Colour
Crystal Form
Composition
Other
Description
1. CaCO3
2.
Anhydrite
Gypsum
3. NaCl
4. MgSO4
5. MgCl
6.
Bitterns:
7. KCl
b. Repeat Usiglio’s experiment. Obtain a 20 mL sample of natural sea water. Evaporate it until
crystals start to form when about ~98% of the liquid has boiled off. Describe your first crystal
residue and test it with a drop of acid. _______________________________________________
__________________________________________________________________________ (2)
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c. Continue evaporating your beaker of sea water to dryness. Describe the final crystals that form
and try and identify what they are. ________________________________________________
___________________________________________________________________________
___________________________________________________________________________(2)
10. In geological history major periods of evaporation have removed salts from the ocean and
left sedimentary salt deposits now to be found on land. The following is a list of ages of
evaporation and major basins where they occur.
Age of Evaporites
Location
1. Quaternary
Utah, Nevada, Chile (Atacama), Western Australia (Shark Bay)
2. Miocene
Mediterranean, Poland
3. Mesozoic (Cretaceous)
US Gulf Coast, Yucatan,
4. Triassic (& Jurassic)
Atlantic all continents, Nova Scotia, New Brunswick, Bath UK
5. Permian
Zechstein: (Germany, Poland) & West Texas
6. Pennsylvanian
Paradox Basin Utah
7. Mississippian
Michigan, Ontario, Ohio
8. Devonian
Williston North Dakota & Saskatchewan
9. Silurian
Michigan, Ohio, Ontario, New York
a. Plot the number and age on the map below (next page): _________________________ (8)
b. Reflect on how particular deposits may influence the oceanographic conditions, biology and
ecology, ground water quality and mineable resources in some of their respective locations.
Discuss at least one example of each of the above items giving its deposit name, location and
influence. _______________________________________________________________ (8)
Oceanography: ____________________________________________________________
Biology/ecology: ___________________________________________________________
Ground Water Quality: ______________________________________________________
Economic Deposit: _________________________________________________________
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