Oceans contain…

CO₂(g) from the atmosphere dissolves in the ocean and reacts with the water.

Reactions (2) and (3) are shown, for simplicity, as Arrhenius acid-base reactions, instead of Brönsted acid-base reactions. The double-headed arrows indicate that these reactions are equilibria. The concentrations of the chemical species depend on the exact conditions of the system (including other chemical species not shown in the reactions).

Indicator Color Chart

Figure 1 shows that the pH of CO₂ dissolved in pure water is about 5. The pK_a for reaction (2) is about 6.4 (K_a = 4.0 x 10^-7). The equilibrium constant expression for reaction (2) can be written as

             
This expression can be rearranged to find the [HOCO₂–(aq)] to [CO₂(aq)] ratio.

             
This low value for the ratio shows that most of the dissolved CO₂ is present as CO₂(aq).

On the other hand, the usual pH of seawater is about 8.2. At this pH in the oceans, about 90% of the dissolved CO₂ is present as the hydrogen carbonate anion, HOCO₂^-(aq), with most of the rest as carbonate ion, CO₃^2-(aq). The major cations in seawater are Na⁺, K⁺, Mg²⁺, and Ca²⁺. The calcium ion, Ca²⁺(aq), is most important for the reactions involving dissolved carbon dioxide that are examined in these activities. (Conventionally, the hydrogen carbonate anion is often written as HCO₃^-, but this representation disguises the fact that the hydrogen is bonded to oxygen. Almost all the acids in nature have one or more hydrogen atoms bonded to oxygen atoms.)

A. Reaction(s) of seawater-dissolved CO₂ with Ca²⁺

Procedures, observations, and analyses

1.

Describe in as much detail as possible (including pH) the aqueous solutions of CaCl₂ and Na(HOCO₂). You may find the Universal acid-base indicator color chart, Figure 1, useful.

Please don safety goggles/glasses and prepare to observe and record all changes that occur. Uncap both test tubes, put the caps in the cup, and pour the CaCl₂ solution into the Na(HOCO₂) solution. Do NOT recap the test tubes (leave the caps in the cup) as you continue to observe the mixture and record what you observe for a minute or two.

Textbooks often suggest that formation of a gas, formation of a precipitate, color changes, or temperature changes are signals that a chemical reaction has occurred in a system.

2.

Was a gas formed in your system? What is your evidence? If gas is formed, what gas (or gases) do you think it might be? Give the reasoning for your choice.

3.

Was a precipitate formed in your system? What is your evidence? If a precipitate formed, what solid (or solids) do you think it might be? Give the reasoning for your choice.

4.

Was there a color change (or changes) in your system? If a color change (or changes) occurred, describe the change(s). If a color change (or changes) occurred, what do you think was going on in the system to cause the change? Give the reasoning for your choice.

5.

Was there a change in temperature in your system? What is your evidence? If there was a change in temperature, an energy change must have occurred. Would this energy change make a positive or negative contribution to the net entropy change, ΔS_net. for the process that occurs? Give the reasoning for your choice.

Discard the reaction mixture down the drain with plenty of running water, thoroughly rinse out the test tubes, and recap them.

Your observations and reasoning in the previous items probably convinced you that a chemical reaction occurred and helped you identify the reaction products, when aqueous CaCl₂ and Na(HOCO₂) solutions were mixed. The Na(HOCO₂) solution represents CO₂ dissolved in seawater, where the major species formed as the gas dissolves is the hydrogen carbonate anion, HOCO₂^-(aq). In order to make the reaction faster and easier to analyze, the concentrations used here are higher than in seawater.

6.

Write and balance a net reaction equation (without spectator ions) describing the reaction when aqueous CaCl₂ and Na(HOCO₂) solutions are mixed. Does this reaction help explain how “CO₂ is also incorporated into the structures of many of these phytoplankton and other organisms”?

B. Reaction(s) of dissolved CO₂ with CaCO₃

The reaction you wrote in item 6 is the reaction that many phytoplankton and other marine organisms use to provide the calcium carbonate, CaCO₃, to build their solid internal and external structures. Now we will further investigate interactions as more carbon dioxide is added to such a system.

Put about 25 mL of limewater (dilute calcium hydroxide, Ca(OH)₂ solution) in a 100-mL beaker or equivalent clear, colorless plastic cup. Add a few drops of Universal acid-base indicator.

7.

What color is the limewater and indicator solution? What does this tell you about the pH of the solution? Is this what you would expect? Explain why or why not.

Please don safety goggles/glasses and prepare to observe what happens when you use a carbon dioxide cartridge and delivery device to add CO₂ gas to the limewater and indicator solution (without splashing all over the place). Add a short, 2-3 second, burst of CO₂ gas to the solution and observe any changes that occur. (If you do not detect any change, add another short burst of gas.)

8.

What change(s) did you observe when CO₂ gas was added to the limewater solution? Is(are) the concentration(s) of any species changing in the solution? What is your evidence? Are new chemical compounds being formed? What is your evidence? Write chemical reaction expressions that explain your observations.

9.

What further change(s) (if any) did you observe? Were any changes surprising? If so, explain why you were surprised. How did you decide when to stop bubbling and that no further change(s) would occur? What was your evidence? Write a chemical reaction expression (or expressions) that explain your observations. Is there any relationship between the reaction(s) here and the reaction characterized by the reaction expression in item 6? Clearly explain your response.

10.

What effect(s), if any, might the kind of change(s) you observed here have on the phytoplankton, corals, and other marine organisms with calcium carbonate structures? Explain the evidence for your answer.

Instructor/presenter notes

The overall objectives of these Activities are to “discover” the seawater reaction between calcium ions and bicarbonate ions (formed by dissolution of carbon dioxide in seawater) and the fate of the resulting calcium carbonate as more carbon dioxide dissolves. The “discoveries” are designed to be based on observations on reacting systems and information from past experience, including other activities that might lead up to these. To begin, in Part A, groups are directed to examine their solutions closely and they should report these before going on. The idea is to get them to really notice the initial characteristics of their solutions (clear and colorless like water and clear and colored with a mildly basic pH), so their changes are as evident as possible. To make sure the focus is only on these solutions, the materials for Part B should not be distributed until Part A is finished.

The next part of the worksheet is more directive than it needs to be. In particular, items 2 through 5 could be condensed into the question, “What observations did your group make?” After groups have carried out the Activity and had five minutes or so to discuss their observations, solicit a single observation from each group and immediately discuss its implications, along the lines suggested in the worksheet items. The idea is for the groups all to agree on the observations made and, as a result of the whole-group discussions, have a common understanding of what has been learned from them. Armed with the results (reactants and probable reaction products), each group can spend five to ten minutes writing a reaction expression, item 6. The groups can then share their reactions and come to a consensus that explains the observations.

Part B begins in the same way with groups looking closely at their reactant solution and inferring its properties. It is worth asking whether (and why) they think carbon dioxide will dissolve and react in this solution, before carrying out addition of the gas. After the reaction(s) are complete, the idea is to end up with an overall reaction that is the reverse of the one written in item 6.

The expected observations and interpretations are given here followed by more general discussion of the implications for ocean acidification, and an extension to other Earth processes, formation of limestone caves and their interiors.

Part A

1.

The clear, colorless CaCl₂ solution looks just like water, so little else can be inferred, except that the ions are colorless. The clear, green Na(HOCO₂) solution obviously has a story to tell. The Universal indicator color chart shows that green signals a pH of about 7 to 8, so a solution containing hydrogen carbonate ion is modestly basic, that is, has a bit more hydroxide ion than hydronium ion.

2.

A gas is formed in the reaction. The evidence consists of bubbles visible in the liquid in the test tube, usually some foaming at the surface of the liquid, and an audible signal, the sound of the bubbles “popping” as they reach the surface. Since this is an activity that is supposed to be about carbon dioxide chemistry in the ocean, it makes some sense to conclude that the gas is probably carbon dioxide, CO₂(g). Carbonates and bicarbonates are also generally known as good sources of carbon dioxide in other reacting systems.

3.

A precipitate is formed in the reaction. The mixture becomes cloudy and opaque as a white solid forms. The only cation present that forms insoluble salts is calcium ion. Calcium carbonate, CaCO₃(s), (chalk, limestone, marble) is insoluble, so it seems a likely candidate for the identity of the solid. (Calcium hydroxide, Ca(OH)₂(s), is also only slightly soluble, but its presence would suggest a basic solution and the next item shows that the solution becomes acidic.)

4.

A color change occurs when the solutions are mixed. The green solution becomes yellow (with a white precipitate as background). The indicator color chart shows that yellow corresponds to a pH of about 6, more acidic than the original bicarbonate solution. This makes sense, if the solution now contains dissolved carbon dioxide, CO₂(aq), (evidenced by the formation of gas in the reaction, item 2), which would make a weakly acidic solution.

5.

Some students holding the bicarbonate test tubes when the solutions are mixed report that the test tube gets a bit cooler. This would mean an endothermic reaction has occurred. That is, the ΔE (ΔH) > 0 for the reaction occurring in the test tube. The temperature change is not large and not everyone observes it, but usually enough groups do report the cooling to make it reasonable to accept this result.

The relationship among entropy, energy (enthalpy), and temperature is the subject of Entropy analyses: Reaction energies and entropies. [This Activity is in development.] If your students have not done this activity or do not have a similar background with the subject, you should omit any analysis based on entropy in this activity. The relationship, ΔS_net = – ΔH_syst/T + ΔS_syst, is useful for the analysis here. An endothermic reaction, ΔH > 0, makes a negative contribution to the net entropy change, ΔS_net, for the reaction and makes it less favorable. Since the temperature change is so small, the enthalpy change must be small as well, so it will have little effect on ΔS_net. There are two much larger positive contributions from ΔS_syst, formation of a gas and formation of the precipitate. Production of a gaseous product always makes a positive contribution to ΔS_syst for the reaction in question and is a favorable effect. The formation of an insoluble salt, like CaCO₃(s) in this reaction, also makes a positive contribution to ΔS_syst. The reason for this is explored in Solubility patterns for ionic compounds. [This Activity is in development.]

6.

The Cl^- and Na⁺ ions do not enter into any of the possible products discussed above, so they can be ignored as spectator ions in writing a net ionic equation. The unbalanced skeleton reaction, using dissolved carbon dioxide, CO₂(aq), as a product, is:

       Ca²⁺(aq) + HOCO₂^-(aq) → CO₂(aq) + CaCO₃(s)

This is balanced in no element except Ca and, in particular, needs another reactant C:

       Ca²⁺(aq) + 2 HOCO₂^-(aq) → CO₂(aq) + CaCO₃(s)

Now Ca, C, and charge are balanced, but H and O are not. Adding H₂O does the trick.

Part B

7.

The Universal indicator color in limewater is blue. The solution is quite basic, pH about 12, as expected for a solution of a hydroxide, calcium hydroxide, even though the salt is fairly insoluble and rather dilute.

8.

Addition of carbon dioxide to the limewater solution produces a white precipitate (a reaction that is often used as a way to test whether a gas is carbon dioxide) and a color change of the indicator toward less basic conditions. “Less basic” means that hydroxide ions must have been used up by the acidic reactions of dissolving carbon dioxide. In the initially quite basic limewater solution, dissolving carbon dioxide reacts rapidly with hydroxide ion.

And the hydrogen carbonate ion rapidly reacts further with hydroxide ion.

It makes sense that the carbonate ions formed then react with calcium ions to form the white calcium carbonate precipitate.

9.

As carbon dioxide continues to be bubbled into the less basic solution, the color changes to yellow, mildly acidic, as the concentration of dissolved carbon dioxide increases. The precipitate also begins to disappear and finally vanishes completely to leave a clear yellow solution. These are the reactions responsible for the observations.

The calcium carbonate dissolves, reaction (9), as carbonate ion is used up, reaction (8), by reaction with the hydronium ion produced by the dissolving carbon dioxide. The sum of these three individual reactions is the overall change in the solution:

Overall reaction (10) is exactly the reverse of reaction (4) written in item 6 to explain the observations upon mixing solutions of calcium ions and bicarbonate ions. Thus, adding carbon dioxide to the system containing calcium carbonate results in dissolution of the solid by driving the reaction in reverse (LeChatelier’s principle) as the carbonate ion concentration decreases, reaction (8). (For a more extensive discussion of the reaction of dissolved carbon dioxide at higher and lower hydroxide ion concentrations, see Aqueous carbon dioxide solutions: Equilibria and rates.)

The sawtooth curve in panel A of Figure 2 is the Keeling curve, the concentration of CO₂ in the atmosphere. As the amount of CO₂ in the atmosphere increases (due largely to fossil fuel burning), the purple squares show that the amount dissolved in the oceans increases in parallel. This is an example of Henry’s law, which says that the concentration of a solute gas increases in direct proportion to the pressure of the gas over the solution. As the concentration of dissolved CO₂ increases, panel B shows that the pH of seawater decreases. This makes CaCO₃(s) more soluble, as the balance between the concentrations of carbonate and hydrogen carbonate ions shifts away from carbonate, reaction (8).

This consequence of this shift is shown in panel C, which plots, Ω_aragonite, the supersaturation of the aragonite crystalline form of calcium carbonate as the pH decreases. This supersaturation is defined as

                

where the concentrations in the numerator are those measured in the seawater and the solubility product in the denominator is the thermodynamic equilibrium value. Figure 2, panel C, shows that Ω_aragonite > 1, so the ion product is larger than the solubility product. That is, the seawater is supersaturated with respect to the solubility of solid calcium carbonate in its aragonite form. Organisms (like phytoplankton, corals, mollusks, etc.) that use calcium carbonate as part of their structures evolved to depend on this supersaturation for their structure building. As more CO₂ dissolves and the concentration of  CO₃^2-(aq) declines, the supersaturation, Ω_aragonite, decreases, Figure 2, panel C, and these organisms may find it more difficult to build needed structures.

10.

Another connection to Earth systems and climate

The reaction (10) in item 9 produces a solution containing calcium ions and bicarbonate ions (hydrogen carbonate ions). This reaction often occurs underground when rainwater saturated with CO₂ seeps through a layer of limestone. As the water dissolves calcium carbonate, it forms openings in the limestone. Caves from which the limestone has been dissolved are often prevalent in areas where there are large deposits of CaCO₃ (for example, Mammoth Cave, Carlsbad Caverns, and Cave of the Mounds in the United States). If the water containing dissolved Ca²⁺(aq) and HOCO₂^-(aq) reaches the ceiling of a cavern, the water will evaporate. As it evaporates, the concentrations of the ions increases and the reactions above begin to run in reverse — CO₂(g) escapes and CaCO₃(s) deposits on the ceiling. Over tens of thousands to millions of years, this deposit can grow into a stalactite (speleothem), a limestone “icicle-like” structure hanging from the ceiling, as shown in Figure 3. If the solution drips to the floor of the cavern and evaporates there, it can build up into a stalagmite (sort of the mirror image of a stalactite). Given enough time, a stalactite and stalagmite below it can grow together to form a column.

The formation of the speleotherms is seasonal and creates annual layers, rather like tree rings, that vary in thickness and composition (isotopic ratios and other minerals that were dissolved along the way from the surface). The layers carry information from the surface that is climate dependent (temperature, rainfall, etc.) and has been used in climate reconstructions that can extend back hundreds to thousands of years and help us learn about today’s changes based on past history.

This cross section of a stalagmite reveals a sequence of layers, laid down over time. Researchers determine the age of the rings using Uranium-Thorium radioisotopic dating, and examine ring thickness and oxygen isotopes to determine past climate. (Photograph copyright Paul Williams, New Zealand National Institute of Water & Atmospheric Research.) From NASA Earth Observatory, Holli Riebeek, “Paleoclimatology: Written in the Earth,” 28 June 2005.

Reference

American Chemical Society Climate Science Toolkit. Go to “Oceans, Ice & Rocks” > “Ocean Chemistry”.

(Adjust volumes for the number of groups doing the Activity.)

Calcium chloride, CaCl₂, solution. Dissolve 12 g CaCl₂ in 100 mL water. Allow the solution to cool before using.

Sodium hydrogen carbonate, Na(HOCO₂), solution. Dissolve 10 g Na(HOCO₂) in 100 mL water. Add 1 mL acid-base indicator and mix.

Acid-base indicator solution. Universal indicator is best, but bromothymol blue or other indicator that changes color around pH 6 is satisfactory, but not quite as informative.

Limewater solution. Add a heaping teaspoonful of calcium hydroxide, Ca(OH)₂, to about a liter of water and stir (magnetically) overnight in a tightly stoppered flask or bottle. Let the mixture settle for several hours and then carefully decant half to two-thirds of the supernatant solution through fine filter paper into a container that can be tightly sealed to prevent CO₂ from the air getting into the solution. (If the solution is not clear and colorless, filter it again.) Shortly before using, dilute an appropriate amount of the solution with an equal part of water. (Undiluted solution will work for the activity, but will take a lot longer to clear and require more CO₂.)

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