Shell Shifts

None needed.

Put on safety goggles.

Place a shell in a cup and cover it with vinegar. What do you notice?

Make a sodium bicarbonate solution by adding one spoonful of baking soda to one cup of water. Make a separate calcium chloride solution by adding one spoonful of calcium chloride to one cup of water. Stir each solution well. How would you describe the two different solutions?

In a new cup, pour in equal amounts of the sodium bicarbonate and calcium chloride solutions—mix well. What do you notice when the two solutions mix? (Click to enlarge image below.) How is this different from the original solutions? 

(Optional: At this point, you can add in a small amount of a pH indicator to keep track of the pH throughout the experiment. If you don’t have any, the experiment will still work without it.)

Add small amounts of vinegar to the cup with the mixed solution until you notice a change. Keep the solution well mixed by stirring.

Add in small amounts of the 0.25M sodium hydroxide to the mixed solution until you notice a change.

The primary component of most seashells is calcium carbonate (CaCO₃). You may have noticed bubbles forming when you initially covered the seashell with vinegar. The bubbles are carbon dioxide (CO₂), which is created when the CaCO₃ in the shell is exposed to an acid such as vinegar. The reaction that describes this process is:

CaCO₃ + 2H⁺ ⇔ Ca²⁺ + CO₂ + H₂O

Sodium bicarbonate (NaHCO₃) and calcium chloride (CaCl₂) both dissolve well in water, and the solutions you made with them should have appeared relatively clear. These dissolutions can be expressed as:

NaHCO₃ → Na⁺ + HCO₃⁻
CaCl₂ → Ca²⁺ + 2Cl⁻

In water, bicarbonate (HCO₃⁻) is never present by itself, but exists in equilibrium with other forms of dissolved inorganic carbon: carbonic acid (H₂CO₃), carbonate (CO₃²⁻), and carbon dioxide (CO₂). The balance between all of these species is shown by these equations:

CO₂ + H₂O ⇔ H₂CO₃ ⇔ HCO₃⁻ + H⁺ ⇔ CO₃²⁻ + 2H⁺

When you mix the NaHCO₃ and CaCl₂ solutions together, the carbonate ion (CO₃²⁻) reacts with the calcium ion (Ca²⁺) to form CaCO₃ through the following reaction:

CO₃²⁻  + Ca²⁺ ⇔ CaCO₃

The mixture should turn cloudy since CaCO₃ is not very soluble in water and will precipitate out of the solution upon forming. You have just made little bits of shells!

When you add vinegar to this mixture, the excess hydrogen ions (H⁺) will dissolve the calcium carbonate (CaCO₃) particles—just like it did to your seashell—and the solution should turn clear again. There should be many bubbles because the bicarbonate in solution will also react with the vinegar to form CO₂. 

Adding a base such as sodium hydroxide (NaOH) will then shift the equilibrium back to solid CaCO₃, forming a cloudy precipitate again. This shows how sensitive CaCO₃ is to the pH of its environment.

A wide variety of ocean organisms—from shellfish and corals to certain kinds of algae—contain calcium carbonate in their exoskeletons. There needs to be a sufficient concentration of carbonate ions available for these creatures to construct their shells. 

Increasing levels of CO₂ in the atmosphere are creating an increase in levels of dissolved inorganic carbon and a decrease of the pH in the oceans, a phenomenon called ocean acidification. 

The carbon species you worked with in this activity are all in a dynamic equilibrium described by these equations:

CO₂ + H₂O ⇔ H₂CO₃ ⇔ HCO₃⁻ + H⁺ ⇔ CO₃²⁻ + 2H⁺

The equilibrium ratios of different carbon species in seawater depend on the pH, shown by the graph below (click to enlarge). The concentration of carbon acid (H₂CO₃) is low, so it has been left off the graph. At ocean pH (shown by the vertical blue bar), the bicarbonate form (HCO₃⁻) represents over 90% of the dissolved inorganic carbon. 

The minerals in the ocean contain large amounts of carbonate, so another reaction that occurs is:

CO₂ + H₂O + CO₃²⁻⇔ 2HCO₃⁻

These two reactions show how increasing CO₂ can lower the pH and reduce the concentration of available carbonate ion. The lower concentration of carbonate reduces the amount available for calcifying organisms that span the food chain and also drives the equilibrium to dissolve more CaCO₃ rocks in the ocean.