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. 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)
(Eqn.1)
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. This is the first part of the rock cycle for chemical weathering.
Because of this weak 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 and clays or oxide/hydroxide residues.
Even Quartz SiO2 can dissolve this way as follows:
SiO2 (s) + 2 H+ +2 e- → H4SiO4 (aq)
When solutions are cold and acidic, there is lots of dissolved 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 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)
(Eqn.2)
Other common ions like Na+ or Mg+2 are too soluble to precipitate until 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 places, in hot springs or in evaporating salty shallow seas or in playa lakes
(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 water or other sediment and kept away from acid
rain, they can stay there and serve to reduce the environmental load of CO2 in the air. This kind of
natural carbon storage will eventually reduce the CO2 in the air but, left to its own slow pace, this will
takes tens of millions of years. Meanwhile we are making more by burning more fossil fuels and slash in
tropical rain forests. We also heat up crushed limestone to get Ca+2 for Portland cement and release even
more CO2 this way!
Explain how weathering of rocks and the precipitation of carbonate minerals is Earth’s long term natural
system for reducing CO2 as a green house gas. (4 points)
1
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
Measure out 10 mL of calcium chloride, CaCl2 solution in a graduated cylinder. This provides the Ca2+
ions as in equation (2) above. In a separate graduated cylinder, measure out 10 mL of sodium carbonate,
Na2CO3 solution. This provides the CO32- ions as in equation (2) above. Observe the nature of the
solutions and describe them here.
(2)
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!
Pour both solutions into the same small beaker. Record your observations.
(2)
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 freezing.
CaCl2 (aq) + Na2CO3 (aq)  CaCO3 (s) + 2NaCl (aq)
(3)
It is possible to separate the calcium carbonate from the sodium chloride in solution in 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 filter paper. Filter paper ________________ g
(1)
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).
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)
While your filter paper and precipitate are drying out, evaporate your little beaker of <20 mL of
solution. According to reaction 3 above, this should be just salt water. Once it dries out and cools off,
examine the crystal residue under a hand lens or binocular microscope and draw and describe the crystal
shapes you see.
(2)
2
Obtain a separate sample of natural calcium carbonate 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)
(4)
The carbonic acid then does equation 1 backwards:
H2CO3 (aq) → CO2 (g) + H2O (l)
(1R)
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 natural waters.
(2)
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. Strangely enough, NaCl is not the
first mineral to form. The order he found was:
Usiglio’s experiment
1. CaCO3
Calcite
2. CaSO4 or CaSO4 – 2H2O
Anhydrite or Gypsum
3. NaCl
Halite
4. MgSO4
Epsomite
5. MgCl
Bischofite
6. Complex K-Mg-Cl-SO4-OH & KCl Bitterns & Sylvite
It is geologically common to find the 1st 3 minerals 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.
Speculate as to why the 1st three are so common while the others are not. What do you think happens to
the water to allow the former ones to form most often but not the latter ones? (4 points)
3
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 in appearance,
water content or hygroscopic nature, solubility etc. Taste for NaCl versus KCl.
Naturally occurring “Evaportie Minerals”
points)
(describe and sketch)
1. CaCO3
Calcite
2. CaSO4 or CaSO4 – 2H2O
Anhydrite or Gypsum
3. NaCl
Halite
4. MgSO4
Epsomite
5. MgCl
Bischofite
6. Complex K-Mg-Cl-SO4-OH Bitterns & KCl Sylvite
4
(12
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
Possible Cause
1. Late Pleistocene
Shark Bay West Australia
2. Miocene
Mediterranean
3. Mesozoic
US Gulf Coast
4. Triassic
Atlantic margins all continents
5. Permian
Zechstein Germany & West Texas
6. Pennsylvanian
Paradox Basin Utah
7. Mississippian
Michigan
8. Devonian
Williston North Dakota & Saskatchewan
9. Silurian
Michigan, Ohio, Ontario, New York
Find these on the world map and on line and think about the different sea level conditions and
environments at those locations when the evaporites formed versus today. Write a possible cause to
explain the evaporates of that age in each location from the following choices: More than one may apply
but put at least one choice for each basin above. (9 points)
A. Cooler Climate with Lower Sea Level to isolate coastal basins
B. Warmer Climate with higher sea level to isolate continental interior basins
C. Shallow rift basin on Continental margin (Red Sea)
D. Horse Latitudes Location or Rainshadow Desert
E. Restricted Basin (cut off from Ocean by faulting or uplift)
F. Tectonically controlled Continental Subsidence making broad shallow inland seas
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 points)
Continue evaporating your beaker of sea water to dryness. Describe the final crystals that form and try
and identify what they are.
(2 points)
5