Carbon Captured Chalk

By Tony Butterfield, based off of a module developed by Alissa Park. Developed by The NSF Research Coordination Network on CCUS (


This module demonstrates several important chemical engineering concepts. Two solutions of soluble salts (Na2CO3and CaCl2) are mixed and undergo a double substitution reaction. The resulting product is table salt (NaCl) and chalk (calcium carbonate). Concepts of solubility, enthalpy of solutions, reaction kinetics are considered. A very similar process could be used by chemical engineers to help relieve the burden of carbon dioxide on our atmosphere by trapping our emissions in benign products, such as chalk. IMPORTANT!!! Under no circumstances should an unsupervised minor perform the procedures described herein. All the following described experiments and methods should be supervised by an adult who is completely familiar with and takes full responsibility for all possible hazards.

General Information

Main Curriculum Tie:
Additional Curriculum Ties:
Environmental Science,Separations
Career Connections:
With growing concern over global climate change, many industries and governmental bodies have become interested in development of technologies that may be used to keep greenhouse gases out of our atmosphere. Chemical, environmental, and civil engineers play a critical role solving climate problems, and this module gives a hands-on demonstration of one such technology.
Mean Time Frame:
30 min
Group Size:
Teams of 2-3, 50 students max
Student Prior Knowledge:
Basic Chemistry

Essential Questions

  • What is a double replacement reaction?
  • What is solubility?
  • How can we remove carbon dioxide from our atmosphere?


Alissa Park, OSCAR (Ohio State Carbonation and Ash Reactivation) Process, teaching module.

Materials & Methods


  • Sodium Carbonate (aka Na2CO3, soda ash, washing soda). Can be purchased from chemical supplier or even at
  • Calcium Chloride (aka CaCl2). Can again be purchased from a chemical supplier, or more easily as ice melt (
  • Water.
  • Scale, some means to measure weight. We bought a cheap onefrom and it works great (This module seemingly brought to you by…).
  • Weighing boats or weighing paper to keep your scale from getting dirty.
  • Spatula, to finely measure out solids on the scale.
  • 50 ml disposable centrifuge tube, or any container that can hold the volume of liquid needed. The centrifuge tubes are good because they can be sealed and shaken to dissolve solids. A stir plate could also be used with, say, a beaker.
  • Filter paper (we’re using #4 filter paper, but coffee filters can work passably). If you want to filter quickly a vacuum filter could be added but gravity filtration works here too.
  • Safety glasses and gloves. This is a pretty safe activity, but it’s always good to use gloves and glasses in a lab environment.


  1. Place a weighing boat on the scale and tare it out.
  2. Measure out 8g of Na2CO3 and dump it into an empty 50 ml centrifuge tube. You may wish to begin this module with the Carbon Capture in Water module to illustrate how CO2 might get from the atmosphere into a compound such as sodium carbonate.
  3. In the same manner, measure out 8 g of CaCl2. Alternatively, students can calculate exactly how many grams of CaCl2needed to react one-to-one with the Na2CO3, and measure that out, but it will be near 8 g anyway.
  4. Dump the CaCl2 into an empty 50 ml centrifuge tube.
  5. With a third 50 ml centrifuge tube or a graduated cylinder, measure out 20 ml of water and pour that water in with the Na2CO3.
  6. The solution will heat up as the Na2CO3 dissolves. Discuss this phenomena with the students.
  7. To completely dissolve the salt put the lid on the tube and shake the solution, periodically stopping to open the lid to release any pressure that may build up due to the heat.
  8. Measure out the same amount of water and mix it with the CaCl2 in the other centrifuge tube. Mix in the same manner as in Step 7.
  9. You should now have two tubes with two different solutions in them.
  1. Pour the CaCl2 solution into the centrifuge tube containing the Na2CO3 solution. If you pour slowly students can see the difference in solution density and can witness their interaction at the interface between the two.
  2. Put the lid on the tube and mix the two solutions together by shaking. You will notice the solution go opaque as it initially forms a gel. After about 3 min of shaking, the solution should become far less viscous and look a lot like milk as the chalk forms minute particles.
  1. Separate the water from the chalk by folding filter paper into a cone over an empty 50 ml centrifuge tube and pouring your solution into it. The water should be collected in the tube and the chalk will remain in the filter paper. You may also wish to use vacuum filtration or a centrifuge if you have it.
  1. Finally, place the chalk out to dry and by the next day you should have a dried product.

Background for Teachers

The theory behind this module is based on several relatively simple concepts.

Enthalpy of Solution

In this module twice we start with a dry salt and end up with that salt dissolved into solution. Consider what happens to the atoms in CaCl2 during this process. Before they can enter the liquid they have to disassociate from the CaCl2 surrounding them. They must break bonds in their crystalline structure, which may be seen as pits of potential energy that they have to escape; this takes energy. We call the energy needed the lattice dissociation enthalpy.

To get into solution an atom from CaCl2 must next become surrounded by solvent molecules, which is water in our case. The energy required to do this is called the enthalpy of hydration, and it is always negative, meaning that the potential energy of the system is always lower with the salt ions surrounded by water than if they are disassociated in a vacuum. This also means that energy is released in this process.

So we have a process that takes energy (breaking the salt’s lattice), and a process that releases energy (surrounding ions by water). The difference between the two will result in either an energy release or uptake, a temperature increase or decrease. If the enthalpy of hydration is more negative than the lattice dissociation enthalpy is position, then the total enthalpy change will be negative and energy will be released into the system, resulting in a temperature increase. This is the case we have with both our salts, CaCl2 and Na2CO3(though this is not the case for all solutes). As the students dissolve both of them, there will be a notable increase in temperature in the solution.

Double Replacement Reaction & Solubility

While in solution, CaCl2 dissociates into two ions of chlorine and one ion of calcium, while Na2CO3 dissociates into two ions of sodium and one carbonate ion. These ions may interact with each other, and once a calcium ion meets up with a carbonate ion they form calcium carbonate (chalk) and sodium chloride, table salt.

Na2CO3 + CaCl2 → CaCO3 + 2 NaCl

This reaction strongly favors the forward direction. Chalk is not very soluble in water (it’s solubility is about 13 mg/L) as one could easily discover with a piece of chalk and a glass of water. It is takes less energy for calcium chloride to interact with itself than to interact with molecules of water in solution. As such, the concentration of chalk in aqueous form is continually held low as it precipitates out of solution to form a solid which will not participate in the reaction.

Because a product in the reaction is continually being removed from solution, and the rate of reaction in the reverse direction depends on the concentration of calcium carbonate, the reaction shown above is forced to the right.



Intended Learning Outcomes

  • Students will understand the difference between soluble and insoluble.
  • Students will be able to define heats of solutions.
  • Students will be able to identify a means to sequester carbon dioxide.

Instructional Procedures

  1. Discuss the importance of carbon sequestration technology.
  2. Explain how CO2 may be converted into carbonic acid and then sodium carbonate.
  3. Split the class into teams of two or three.
  4. Assign out tasks as described in the Methods section of this module within each group.
  5. Ask students to pay attention to the temperature of each solution as they dissolve their salts.
  6. Ask students to pay attention to and make observations of the viscosity of their mixture of the two solutions.
  7. Finish the methods and, as the chalk filters, ask students to perform the assessment questions in this module.

Optional Activities & Extensions

  • Couple this module with our Carbon Capture in Water module to illustrate how carbon dioxide can be removed from the atmosphere and put into substances such as the sodium carbonate reactant used in this module.
  • Have students put a drop of each reactant solutions on a microscope slide and have them observe the microscopic formation of calcium carbonate crystals.
  • Allow students to dye their solution and then press the resulting chalk into a solid piece of chalk as it dries.