1 / 11

THE GEOCHEMISTRY OF NATURAL WATERS

2. LEARNING OBJECTIVES. Understand sources of CO2 in natural waters.Define and understand alkalinity.Learn to calculate the solubility of carbonate minerals such as calcite.Understand the common-ion effect.Become familiar with the concept of incongruent dissolution.Apply these concepts to some

wood
Download Presentation

THE GEOCHEMISTRY OF NATURAL WATERS

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

    1. 1 THE GEOCHEMISTRY OF NATURAL WATERS THE CARBONATE SYSTEM CHAPTER 3 - Kehew (2001) The common-ion effect and incongruent dissolution

    2. 2 LEARNING OBJECTIVES Understand sources of CO2 in natural waters. Define and understand alkalinity. Learn to calculate the solubility of carbonate minerals such as calcite. Understand the common-ion effect. Become familiar with the concept of incongruent dissolution. Apply these concepts to some case studies. In this lecture we will pick up where we left off in Lecture 3, and continue with Chapter 3 in Kehew (2001). First we will consider sources of CO2, other than atmospheric, in natural waters. We will then define the concept of alkalinity and attempt to understand its relevance. The bulk of this lecture is devoted to understanding solubility equilibria of carbonate minerals, such as calcite. This includes being able to calculate the solubility of carbonates under specified conditions of pCO2 and pH, considering the common-ion effect and its relevance to natural waters, and familiarizing ourselves with the concept of incongruent dissolution. With these tools under your belt, you are then directed to the latter half of Chapter 3 in Kehew (2001), where he applies these concepts in a discussion of some interesting case studies. Because Kehew (2001) does a very good job describing these case studies, and we have limited time together, I will not be reviewing the case studies in class. However, I direct you to the Lecture 4 page on the course web site for some study questions to help guide your reading in the text. In this lecture we will pick up where we left off in Lecture 3, and continue with Chapter 3 in Kehew (2001). First we will consider sources of CO2, other than atmospheric, in natural waters. We will then define the concept of alkalinity and attempt to understand its relevance. The bulk of this lecture is devoted to understanding solubility equilibria of carbonate minerals, such as calcite. This includes being able to calculate the solubility of carbonates under specified conditions of pCO2 and pH, considering the common-ion effect and its relevance to natural waters, and familiarizing ourselves with the concept of incongruent dissolution. With these tools under your belt, you are then directed to the latter half of Chapter 3 in Kehew (2001), where he applies these concepts in a discussion of some interesting case studies. Because Kehew (2001) does a very good job describing these case studies, and we have limited time together, I will not be reviewing the case studies in class. However, I direct you to the Lecture 4 page on the course web site for some study questions to help guide your reading in the text.

    3. 3 THE COMMON-ION EFFECT - I Calcite solubility is governed by the reaction: CaCO3(s) ? Ca2+ + CO32- (1) Suppose we added a second compound containing carbonate, and this compound is more soluble than calcite, e.g., Na2CO3. This compound will dissolve according to: Na2CO3(s) ? 2Na+ + CO32- (2) To the extent that reaction (2) proceeds to the right, by Le Chatlier’s principle, this will force reaction (1) to the left, precipitating calcite. According to Le Chatlier’s principle, any increase in the concentration of CO32- will shift the following reaction to the left and cause calcite to precipitate: CaCO3(s) ? Ca2+ + CO32- One way to increase the concentration of CO32- is to add a salt to the solution that both contains carbonate ion, and is more soluble than calcite. For example, Na2CO3 is much more soluble than calcite. Thus, when we add this salt to a solution saturated with calcite, its addition will cause CO32- to increase and therefore, calcite will precipitate. This is called the common-ion effect. The anion CO32- is the common-ion between CaCO3 and Na2CO3. According to Le Chatlier’s principle, any increase in the concentration of CO32- will shift the following reaction to the left and cause calcite to precipitate: CaCO3(s) ? Ca2+ + CO32- One way to increase the concentration of CO32- is to add a salt to the solution that both contains carbonate ion, and is more soluble than calcite. For example, Na2CO3 is much more soluble than calcite. Thus, when we add this salt to a solution saturated with calcite, its addition will cause CO32- to increase and therefore, calcite will precipitate. This is called the common-ion effect. The anion CO32- is the common-ion between CaCO3 and Na2CO3.

    4. 4 THE COMMON-ION EFFECT - II The effect of adding sodium carbonate to the solution can be demonstrated by adjusting the charge-balance expression to be: By repeating the derivation of the equations on a previous slide using this charge-balance expression we obtain: Increasing Na+ concentration leads to decreased Ca2+ concentration.

More Related