This summer, I was fortunate to have spent two weeks participating in research at South Dakota State’s Ice Core and Environmental Chemistry Lab under the guidance of Dr. Cole-Dai. During my career, I always had a vague sense of the functional value ice cores bring to those who study Earth’s history. However, Dr. Cole-Dai and his research group quickly made it evident to me that ice core chemistry involves so much more than what I had considered. I began to understand just how precious these ice cores are. Not simply due to their inherent fragility, but the extent to which they play a crucial role in the accumulation of evidence used to understand Earth’s climate history and its future. Since this two-week stay was part of our master’s program at SDSU, teachers were expected to develop a laboratory activity that reflected aspects of whatever research lab we happened to be working in. Given that the Cole-Dai research team focuses on measuring trace chemical impurities in ice cores, my lab partner and I thought it might be fun to have students simulate the process used to detect past volcanic eruptions by measuring sulfate concentrations at various depths within the “ice core”. This lab provides a great opportunity for students to use different units of concentration beyond molarity, explore how knowledge of chemistry can uncover past events on Earth and practice the skill of generating a calibration curve.
Figure 1: SO2 gas goes through an oxidation process turning into SO42-
To understand the function of this lab, you need to have a basic understanding of the contents within an ice core. Many materials can be studied within a small sample of ice, this includes cations and ions. Among these anions is sulfate, SO42-. This ion is always present in the ice record, but when its concentration is significantly higher for a short period, it can indicate that a significant volcanic eruption took place somewhere on Earth about 1-2 years earlier. By determining how deep in the ice core the sulfate ion concentration peak is found, we can estimate how many years ago an eruption took place. Those volcanic eruptions can be detected in ice cores from all over the ice sheet at the same depth—that is, at the same point in Earth’s history. Though many other ions in the ice samples can tell us other important things about Earth’s past, this places the focus solely on sulfate and it’s assumed that sulfate is the only ion present in the ice core sample for each group.
Figure 2: Actual ion chromatography analysis from ice core indicating mystery volcano in 1458 (Cole-Dai research team)
Prepping the “Ice Core” Samples
With the relationship between high sulfate concentration and volcanic eruptions established, my partner and I figured we could easily have students simulate this dating of volcanic eruptions process by using ice cubes as a proxy for discrete samples of an ice core (Figure 3). Like the samples cut from a core, each ice cube will correspond to a certain depth. Additionally, since sulfate is always present in a core, a small and consistent amount of a stock solution of Na2SO4 is added to each ice cube. However, one of the ice cubes will be given a noticeably greater amount of Na2SO4 and will therefore contain a higher concentration of sulfate. Once each ice tray contains the appropriate solutions of water and Na2SO4, they can then be placed in the freezer until you are ready for the lab. The amount of ice cubes, labeled depths, and volume of Na2SO4 used in this preparation process can easily vary to fit your needs but the exact preparation process we used is described in the teacher notes within the included lab activity below.
Figure 3: Example ice cube tray layout for samples at various depths
Prepping the Standard Solutions
When detecting trace impurities from actual ice cores in the research lab, one of the things I thought was pretty cool was the reliance upon an accurate calibration curve. By using standard solutions of known concentration, the ion chromatograph can measure the conductivity of each solution, resulting in a linear fit. Then, when the conductivity of the discrete sample from the ice core is measured, its concentration can be calculated based on the linear equation derived from the calibration curve. Seeing actual research on ice cores apply a concept that we’ve done in class before sparked an interesting connection between what we do in class and the skills scientists apply in the field. Though students won’t be using an ion chromatograph for this lab, many classrooms have access to sensors that can measure the conductivity of a solution. Therefore, we should be able to reasonably replicate the process described above using a conductivity sensor, standard solutions, and our ice cubes containing sulfate.
Any soluble salt can be used for this lab, so feel free to use whatever you have on hand. Simply prepare a solution that gives you a reasonable conductivity as measured by the probe or meter you have available, and then dilute the solution to give whatever series of standards you’d like. We used Vernier Conductivity sensors for our preparation. For instance, our stock solution (5000 ppm sulfate) had a conductivity of about 8000 microsiemens (µS).
We prepared a 5000 parts per million sulfate (5000 ppm SO42-) solution and then diluted this stock to the required concentrations for our standards. We also used the 5000 ppm stock solution to prepare our group samples in the ice cube trays.
To prepare 5000 ppm sulfate solution from solid Na2SO4, there are two common options:
- From anhydrous Na2SO4, dissolve 7.39 g solid into enough distilled water to yield a final volume of 1.00 L in a volumetric flask or other precise container.
- From sodium sulfate decahydrate (Na2SO4 ∙ 10H2O), dissolve 16.77 g Na2SO4 ∙ 10H2O into enough distilled water to yield a final volume of 1.00 L in a volumetric flask or other precise container.
(Alternatively, you can prepare a 1000 ppm SO42- stock solution (which is also the highest concentration among the standards used for the standardization curve) and dilute that solution to prepare the other standards using one-fifth the amounts listed above for 1.00 L of solution.)
We prepared our standards at 1000 ppm, 500 ppm, 100 ppm, and 25 ppm. Volumes of 100 mL were prepared for each standard using volumetric flasks; another appropriately precise container could be used. Students took small amounts of each standard solution for measurement to establish their own calibration curves; if you’re doing this lab with several sections, a larger amount of each standard solution may be necessary.
Figure 4: Stock and standard solutions setup
Once students have generated their calibration curve by plotting conductivity (µS) against the concentration of the standard solutions, they will determine their line of best fit. If done correctly, this should produce a linear relationship and therefore an equation in the form of y = mx + b. Since X represents concentration, students can then rearrange their empirical equation to solve for X. When it comes to time to measure the conductivity of their “ice core” samples, they can then plug in their conductivity for the Y variable and solve for concentration (X).
Table 1: Example data table from student handout
Like the graph in Figure 2, we thought it would be useful for students to plot their sulfate concentrations against the corresponding depths. This provides a nice visual representation of the data and makes the spike in sulfate obvious. Once the depth at which the spike in sulfate has been determined, students can now begin the process of uncovering which volcano erupted in the past.
Figure 5: Example graph from student handout
Determining which volcano erupted
At this point, students have determined the depth at which a large spike in sulfate concentration has occurred. For this depth to correlate with one of the selected historical volcanic eruptions in the table on the right, there needs to be some way to convert depth into years. While the actual process involved for doing this is quite meticulous, my partner and I went with the following conversion for students to use:
100 years of time = 26.6 meters of ice
Table 2: Possible volcanic eruptions for students to identify within student handout
Once their conversion is made, students can then attempt to identify their volcanic eruption. One more factor to consider is that the ice sheets in Greenland typically aren’t able to show evidence of eruptions for the entire globe. The emissions from volcanoes that are too far south on Earth’s surface don’t reach the far northern latitude of Greenland. To appear in the ice cores collected at Summit, the volcano should be found no further south than around 30°S latitude. If students determine that there are multiple volcanic events that correspond to the sulfate peak at a specific depth in their ice core, they may be able to select the most likely eruption based on latitude. This creates a potential opportunity for students to research the precise location of their volcano and plot it on a global map like the one in figure 6.
Figure 6: Global map used to help students recognize volcano location (world map from printablee.com; modified using PowerPoint)
Though I have yet to implement this lab into my class, I fully intend to find a good fit for it this year. While on campus at SDSU, we were able to run this lab with other teachers who provided us with some great feedback. Overall, the lab went smoothly, and the results turned out great. I really like the idea of using this lab as a form of practicum and evaluating students based on their ability to determine the correct volcanic eruption that had been previously assigned to their group. There are so many fascinating things to uncover when it comes to ice core chemistry and I hope this activity might serve as a useful opportunity to expose your students to some of the awesome mysteries ice cores can help us solve.
A special thanks to Ryan Bleth, a chemistry teacher in Bismarck, North Dakota, for all his work while collaborating with me on this activity.
