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ACTIVITY in Le Châtelier's Principle, Equilibrium Constants. Last updated October 13, 2021.
In this activity, students will use the simulation Surface Ocean pH Levels, from the Chemistry in Context Simulation Suite, to investigate the values related to ocean acidification and interpret systems with multiple equilibria.
High School (AP Chemistry)
AP Chemistry Alignment
This activity will help prepare your students to meet the following learning objectives:
- Unit 7: Equilibrium
- Topic 7.8: Representations of Equilibrium
- TRA-7.F: Represent a system undergoing a reversible reaction with a particulate model.
- Topic 7.9: Introduction to Le Châtelier’s Principle
- TRA-8.A: Identify the response of a system at equilibrium to an external stress, using Le Châtelier’s principle.
- Topic 7.13: pH and Solubility
- SPQ-5.C: Identify the qualitative effect of changes in pH on the solubility of a salt.
- Topic 7.8: Representations of Equilibrium
- AP Chemistry Science Practices
- Practice 4: Model Analysis
- Practice 6: Argumentation
This activity will help prepare your students to meet the performance expectations in the following standards:
- HS-PS1-6: Refine the design of a chemical system by specifying a change in conditions that would produce increased amounts of products at equilibrium.
- Scientific and Engineering Practices:
- Using Mathematics and Computational Thinking
- Developing and Using Models
- Analyzing and Interpreting Data
- Constructing Explanations and Designing Solutions
By the end of this activity, students should be able to:
- Use equilibrium principles to justify the relative proportions of various species in the carbonate buffer system.
- Predict the direction of change for various reactions in simultaneous equilibrium.
- Explain the chemistry involved in ocean acidification due to carbon dioxide emissions.
This activity supports students’ understanding of:
- Equilibrium Reactions
- Equilibrium Constants
- Le Châtelier’s Principle
- Acids and Bases
- Ocean Acidification
Teacher Preparation: minimal
Lesson: 45-60 minutes (depending on prior knowledge)
- Chemistry in Context Simulation: Surface Ocean pH Levels
- URL: https://acswebcontent.acs.org/ChemistryInContextSuite/applets/OceanAcidification/oceanAcid.html
- No specific safety considerations are necessary for this activity.
- Definition of Bronsted-Lowry acids and bases
- Identification of conjugate acid-base pairs
- Acid/base dissociation equations (this activity uses H3O+, rather than H+)
- Simple Ka or Kb expressions
- pH as a function of [H3O+]
- Henry’s Law is not specifically included in this activity, but would be an applicable addition, if desired.
- The simulation includes background notes that are appropriate for the teacher but may be a little complex for students if just learning about systems with multiple equilibria.
- Students should read and complete questions 1-2.
- Review answers and review simple acid/base reactions and identifications.
- Open the simulation and point out the slider and that it shows atmospheric pressure and ocean pH.
- Students read, interact with the simulation, and complete questions 3-4.
- Review answers and review any equilibrium ideas needed, based on student responses.
- Students interact with simulation and complete questions 5-8.
- Review answers and review any qualitative equilibrium concepts that arise.
For the Student
It is a known issue that, in an attempt to meet our energy needs, humans have put a lot of pollution into Earth and its global systems. There have been many changes and much research over recent years to clean up the pollution we’ve already created and to find alternative methods that are less polluting. In this activity, students will investigate three major pollutants that result from the burning of fossil fuels to generate electricity. These are CO2, SO2, and NOx. The “x” in NOx is a scientist’s shorthand to designate any of the group of binary nitrogen-oxygen compounds, such as N2O, NO, and NO2.
Acid rain occurs when certain nonmetal oxides in the atmosphere dissolve and react within water while still in the atmosphere. These nonmetal oxides can react with small amounts of moisture and stay in the air for a while, or they can dissolve in the rain passing through the atmosphere to reach Earth as acid rain. You may recall from your study of acids and bases that metal oxides can react with water to form bases and nonmetal oxides can react with water to form acids.
- Both SO2 and NOx undergo some chemical transformations in the atmosphere before they react with water. The new species are SO3(g) and NO2(g), respectively.
- For each of these new molecules, and for CO2(g), write a balanced chemical equation that shows how each forms an acid upon reaction with water. Include symbols to show the states.
- Write the proper chemical name for each of the three acids.
- One of the three acids is a monoprotic strong acid. Identify this acid, then calculate the pH of a 0.015-M solution of this acid.
- Sulfuric acid and carbonic acid are both diprotic acids. This means they will undergo acid dissociation in water in two steps.
- For sulfuric acid, write the acid dissociation equation for each step.
- Identify the K value for each of the steps you wrote above as being “very high” (K>>1) or “very low” (K<<1) and give a brief explanation for your answers.
- In the beaker below, draw the relative amounts of all species that would be present in a 0.10-M H2SO4 solution.
- Assume a total of 10 H2SO4 molecules prior to any ionization
- You do not need to show any H2O molecules
- Your image will not be quantitatively accurate but should be qualitatively accurate
Of the three acids discussed, carbonic acid is the weakest, and thus has the least impact on acid rain, even though CO2 is more abundant in the atmosphere than the other two pollutants. To understand why carbon dioxide can have an effect on the ocean’s pH, we must look beyond acid rain and consider how the atmosphere interacts with the ocean.
The ocean is a vast and variable mixture of many substances. Though it is not a true solution due to the many large and tiny particles floating within it, there are still a lot of substances that dissolve in the ocean. For this activity, treat the ocean like it is a typical aqueous solution, where water is the solvent and there are many different solutes. Ignore any undissolved substances except when they contribute to relevant dissolved particle amounts through equilibrium factors.
The temperature of the ocean varies, both around the globe and from day to day in any one area. For this activity, we will assume a constant and standard temperature while analyzing the system.
Open the simulation at: https://acswebcontent.acs.org/ChemistryInContextSuite/applets/OceanAcidification/oceanAcid.html
- Use the slider in the simulation to find the pH for every 100 ppm CO2, for the range of concentration shown on the scale. In the space below, sketch a plot of CO2 concentration vs ocean pH.
- Is the relationship linear?
- Describe the pattern that you see, in terms of the plotted variables.
- According to the data, what effect does carbon dioxide have on the acidity of the ocean?
- Consider the equation you wrote in question #1 for how carbon dioxide reacts with water. This reaction can actually be considered as happening in two different steps, each with their own equilibrium positions.
- Write the chemical equation that represents the equilibrium between atmospheric carbon dioxide (CO2(g)), and dissolved carbon dioxide (CO2(aq)).
- Write the chemical equation that represents the equilibrium between dissolved carbon dioxide and carbonic acid.
- What species exists in both equilibrium equations?
- Use equilibrium concepts to explain why there is a correlation between the amount of carbon dioxide in the atmosphere and the amount of carbonic acid in the ocean.
Click the button that says, “Show Graph”. When the graph appears, click the button that says, “Carbonates in a Closed System.” To imagine our ocean equilibria in a closed system, pretend you could box in a specific amount of ocean, so that nothing enters or escapes, and nothing else interacts with the system. Assume that equilibrium levels related to the atmosphere had already been established, and there are no additional interactions with the atmosphere.
- You can see that this graph shows three different species as part of the carbonate system. Using what you know about acid dissociation reactions, write all chemical equations necessary to show how each species forms, starting from the dissolved carbonic acid.
- Use the graph to identify the pH ranges where each of the species is most dominant (pH is the x-axis, and each line represents a different carbonate species):
- Carbonic acid
- Bicarbonate ion
- Carbonate ion
- Use Le Châtelier’s principle to explain why carbonate is dominant in the pH range you specified for this system.
Click the button, “Carbonates in an Open Ocean.” This graph shows a constant concentration of carbonic acid, while the other species vary according to pH. This is different from the closed system because in an open system the carbonic acid is in equilibrium with the carbon dioxide in the atmosphere. When the atmospheric CO2 remains constant, the concentration of carbonic acid remains constant. This graph shows that the proportion between the other two species in the carbonate system varies with pH.
Buffer- A system of chemical species that resist changes in pH because each species is involved in equilibrium reactions that will partially reverse any addition of hydronium or hydroxide.
- Assume that the carbonate system is the major buffering system of the ocean.
- Explain how increasing the amount of CO2 dissolved in the ocean will decrease the amount of the carbonate species in the ocean.
- The shells of many ocean organisms are made primarily of calcium carbonate (CaCO3), which the organisms build from the calcium and carbonate ions in the ocean.
- Write the chemical equation that represents building these shells.
- Use equilibrium principles to explain what will happen to shelled organisms in the ocean if atmospheric CO2 continues to rise.